1. Ch. 20: ACIDS AND BASES Ch. 20: ACIDS AND BASES www.lab-initio.com

2. CA Standards

3. 20.1 Properties of Acids (and Bases)  Acids are proton (hydrogen ion, H + ) donors  Acids have a pH lower than 7  Acids taste sour or tart  Acids affect indicators  Blue litmus turns red Methyl orange  Methyl orange turns red  Acids react with active metals, producing H 2  Acids react with carbonates  Acids neutralize bases Litmus paper (pH indicator) Below pH 4.5Above pH 8.3 Goals of this section: be able to list properties of acids and bases, and know the naming conventions for acids and bases.

4. Acids are Proton (H + ion) Donors Strong acids are assumed to be 100% ionized in solution (good H + donors). Weak acids are usually less than 5% ionized in solution (poor H + donors). HCl Hydrochloric acid H 2 SO 4 sulfuric acid HNO 3 Nitric acid H 3 PO 4 Phosphoric acid CH 3 COOH Acetic acid Organic acids

5. Acids Have a pH less than 7

6. Acids Taste Sour   Citric acid in citrus fruit   Malic acid in sour apples   Lactic acid in sour milk and sore muscles   Butyric acid in rancid butter Organic acids are weak acids. Some are used as flavoring agents in food.

7. Organic Acids Organic acids all contain the “carboxyl” group, COOH, sometimes several of them. CH 3 COOH → CH 3 COO - + H + (acetic acid) The carboxyl group is a poor proton donor, so ALL organic acids are weak acids.

8. Acids Affect Indicators Blue litmus paper turns red in contact with an acid. Methyl orange turns red with addition of an acid

9. Acids React with Active Metals Acids react with active metals to form salts and hydrogen gas. Mg + 2HCl  MgCl 2 + H 2 (g) Zn + 2HCl  ZnCl 2 + H 2 (g) Mg + H 2 SO 4  MgSO 4 + H 2 (g) By the way, “active metals” includes the alkali metals, some alkali earth metals, Cu, Zn and some others. Inactive metals are things like Au and Pt which don’t react, don’t tarnish, etc.

10. Acids React with Carbonates (CO 3 groups) 2 CH 3 COOH + NaHCO 3 acetic acid (vinegar) baking soda CH 3 COONa + H 2 O + CO 2 (g) Sodium acetate water carbon dioxide

11. Effects of Acid Rain on Marble carbonate) Effects of Acid Rain on Marble (calcium carbonate) George Washington: BEFORE George Washington: AFTER

12. Acids Neutralize Bases HCl + NaOH  NaCl + H 2 O Neutralization reactions always produce a salt and water. H 2 SO 4 + 2NaOH  Na 2 SO 4 + 2H 2 O 2HNO 3 + Mg(OH) 2  Mg(NO 3 ) 2 + 2H 2 O In general: HX + YOH → YX + H 2 O acid + base salt + water

13. Properties of Bases  Bases are proton (hydrogen ion, H + ) acceptors  Bases have a pH greater than 7  Bases taste bitter  Bases affect indicators  Red litmus paper turns blue  Phenolphthalein turns purple/pink  Solutions of bases feel slippery  Bases neutralize acids  Bases can be strong or weak electrolytes (conducting electricity)

14. Bases are Proton (H + ion) Acceptors   Sodium hydroxide (lye), NaOH   Potassium hydroxide, KOH   Magnesium hydroxide, Mg(OH) 2   Calcium hydroxide (lime), Ca(OH) 2 OH - (hydroxide) in base combines with H + in acids to form water H + + OH -  H 2 O

15. Bases have a pH greater than 7

16. Bases Affect Indicators Red litmus paper turns blue in contact with a base. Phenolphthalein turns bright pink in a base (it’s orange when in acid).

17. Bases Neutralize Acids Milk of Magnesia contains magnesium hydroxide, Mg(OH) 2, which neutralizes stomach acid, HCl. 2 HCl + Mg(OH) 2 acid base MgCl 2 + 2 H 2 O neutral salt water

18. Names and Formulas of Acids and Bases Acid – a compound that produces hydrogen ions when dissolved in water. General formula for an acid is HX, where X is a monatomic or polyatomic ion. – –Examples: HF, HCl, HNO 3, H 2 SO 4 hydrofluoric, hydrochloric, nitric, sulfuric acids There are three different naming conventions for acids, as seen on the following page. Trivia: Baking powder contains tartaric acid, which is common in grapes and bananas, as well as baking soda (sodium bicarbonate)

19. Names and formulas of Acids and Bases 1. 1.When the name of the anion (X) ends in –ide, the acid name begins with hydro-. The ending turns to –ic and is followed by the word acid. – –Example: HCl(aq) X=chloride: name changes to hydrochloric acid. H 2 S: X=sulfide, name changes to hydrosulfuric acid. 2. 2.When the anion name ends in –ite, the acid name is the stem of the anion with the suffix –ous, followed by the word acid. So H 2 SO 3 (X=sulfite) is named sulfurous acid. 3. 3.When the anion name ends in –ate, the acid name is the stem of the anion with the suffix –ic, followed by the word acid. So HNO 3 (X=nitrate) is named nitric acid.

20. Bases A base is a compound that produces hydroxide ions (OH - ) when dissolved in water. Ionic compounds that are bases are named in the same way as any other ionic compound: – –NaOH sodium hydroxide – –Ca(OH) 2 calcium hydroxide

21. OK, let’s practice naming Sample problem 20-1 Name these compounds: HClO, HCN, H 3 PO 4 – –HClO: the anion, hypochlorite (ClO - ) ends with –ite so change that to –ous, and you get: Hypochlorous acid -HCN: the anion, cyanide, ends in –ide, so change it to –ic and you get: Hydrocyanic acid -H 3 PO 4 : the anion, phosphate, ends in –ate, so change that to –ic and you get: Phosphoric acid

22. More naming practice Name each acid or base: HF, KOH, HNO 3, H 2 SO 4 HF – hydrofluoric acid (from fluoride) KOH – potassium hydroxide (no change it’s a base) HNO 3 – nitric acid (from nitrate) H 2 SO 4 – sulfuric acid (from sulfate) Write formulas for chromic acid (chromate ion is CrO 4 -2 ), iron(II) hydroxide, hydroiodic acid, lithium hydroxide. Chromic acid – H 2 CrO 4 Iron(II) hydroxide – Fe(OH) 2 Hydroiodic acid – HI Lithium hydroxide - LiOH

23. 20.2 Hydrogen ions and acidity Goals of this section: given the H + or OH - concentration, classify a solution as acidic, basic or neutral. Convert hydrogen ion concentrations into values of pH and hydroxide ion concentrations into values of pOH.

24. Self-ionization of water Hydronium ion hydroxide ion H 3 O + OH - A water molecule that loses a hydrogen ion becomes a negatively charged hydroxide ion (OH - ). A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion (H 3 O + ). These ions make water slightly conductive. H+ donor H+ acceptor

25. Hydrogen ions from water The self-ionization of water can be written the following ways: H 2 O(l) ↔ H 3 O + (aq) + OH - (aq) (this way on test!) water hydronium ion hydroxide ion H 2 O(l) ↔ H + (aq) + OH - (aq) water hydrogen ion hydroxide ion In water or aqueous solutions, hydrogen ions (H + ) are always joined to water molecules in the form of hydronium ions (H 3 O + ). Three names they are called are hydrogen ions, hydronium ions or solvated protons. Self-ionization occurs to only a small extent. The concentrations [H + ] = [OH - ] = 1 x 10 -7 moles/liter. If [H + ] = [OH - ] then it’s a neutral solution.

26. Relative concentrations of [H + ] and [OH - ] In any aqueous solution, the concentrations of [H + ] and [OH - ] are interdependent. That means that when [H + ] increases then [OH - ] decreases, and vice versa. Thinking about Le Chatelier’s principle, if additional reactants, [H + ] or [OH - ], are added, that shifts the equilibrium of the solution, in this case, to form more water. Let’s say [H + ] increases: H + (aq) + OH - (aq) ↔ H 2 O(l) Then when it shifts right to makes more water, the concentration [OH - ] goes down because [OH - ] is consumed to make more water.

27. Ion-product constant for water, K w For aqueous solutions, the product of the [H + ] (or [H 3 O + ] - same thing) and the [OH - ] always equals 1x10 -14 M 2. [H + ] x [OH - ] = 1 x 10 -14 M 2 Example: we already talked about the one neutral example where each of [H + ] and [OH - ] =1x10 -7 so in that case, [1x10 -7 ] x [1x10 -7 ] = 1x10 -14 But what about when the solution is NOT neutral? An acidic solution is one where [H + ] > [OH - ], so in that case, [H + ] > 1x10 -7 and therefore, [OH - ] < 1x10 -7 For example, in an acidic solution, you could have [H + ] x [OH - ] = [1x10 -5 ] x [1x10 -9 ] = 1 x 10 -14

28. Ion-product constant for water, K w But what about if the solution is a basic solution instead? (also known as an alkaline solution) A basic solution is one in which the [OH - ] concentration is more than the [H + ] concentration. For example, in a basic solution, you could have [H + ] x [OH - ] = [1x10 -10 ] x [1x10 -4 ] = 1x10 -14 Note that the [OH - ] concentration of 1x10 -4 is greater, than [H + ] at 1x10 -10,but the total product is still 1x10 -14.

29. Sample problem 20-2

30. 20.2 The pH Concept Expressing the hydrogen and hydroxide ion concentrations in exponents of 10 can be cumbersome (as you’ve just seen). A more convenient system to use is the pH system. BUT – it involves math. It involves logarithms. OH NO! Logarithms, the thing that I asked my math teacher if we would ever have any use for! Don’t worry, even if you hate logarithms, you can get the hang of pH, no problem!

32. pH pH was proposed in 1909 by the Danish scientist S Ø ren S ø renson>> The pH scale goes from 0 to 14. Neutral solutions are right in the middle, at 7. Acids go from 0 to 7 on the pH scale. A pH of 0 is strongly acidic. Bases go from 7 to 14 on the pH scale. A pH of 14 is strongly basic. Here is the math related to pH: pH = -log[H + ]

33. pH math review Recall from math that 1. 1.log[a] is saying the same thing as log 10 [a] 2. 2.If y = log 10 [x] then 10 y = x Example: if y = log 10 [100], then 10 y =100 so y must be equal to 2 because 10 2 = 100 3. 3.log [ab] = log [a] + log [b] Example: pH = - log[1x10 -7 ] = -log[1] - log[10 -7 ] Now what is log[1]=y? That means 10 y =1, so y must be equal to zero because 10 0 =1. What is log[10 -7 ]=y? That means 10 y =10 -7, so by inspection y must equal -7. So let’s put it all together: pH = - log[1x10 -7 ] = -log[1] -log[10 -7 ] =-(0+(-7))= 7

34. Logarithm table pH = -log[x] If x (the concentration) equalsThen the pH equals 10 -14 14 10 -12 12 10 -8 8 10 -6 6 10 -4 4 10 -2 2 10 -1 1 10 So even if the logarithm math is hard for you, it’s pretty easy to see the pattern when it comes to pH.

35. pOH The pOH of a solution equals the negative logarithm of the hydroxide ion concentration: pOH =-log [OH - ] The form is similar to the form for pH: pH = -log[H + ] How are those two related? Easy! pH + pOH = 14 Example: If we know that the pH of a solution is 5, what is the pOH? Is that solution an acid or a base? Example: If we know that the pOH of a solution is 13, what is the pH? Is it acid or base?

36. Review: which one is an acid, a base, and neutral? What do you get when you multiple B’s two concentrations together? How about for C?

38. pH + pOH = 14

39. Calculating pH, pOH pH = -log 10 [H 3 O + ] (or –log 10 [H + ]) pOH = -log 10 [OH - ] Relationship between pH and pOH pH + pOH = 14 Finding [H 3 O + ], [OH - ] from pH, pOH [H 3 O + ] = 10 -pH [OH - ] = 10 -pOH SUMMARY PAGE TO PUT ON YOUR STUDY BUDDY!

41. Sample problem 20.3 and 20.4 What is the pH of a solution with a hydrogen ion concentration of 1x10 -10 M? pH = -log(1x10 -10 ) = - ( log(1) + log(10 -10 )) = (0 -(-10)) = 10 The pH of an unknown solution is 6.00. What is its hydrogen ion concentration? -log[H + ] = pH -log 10 [H + ] = 6.00 which says 10 -6 = [H + ] Note that the negative sign had to move from the left to right side of the equation FIRST!

42. Sample problems 20-5

43. Sample problem 20-6 (more challenging) What is [H + ] of a solution if the pH = 3.7? Since pH = -log [H + ] 3.7 = -log [H + ] -3.7 = log 10 [H + ] So 10 -3.7 = [H + ] On your calculator, do 10 ^ -3.7 and you’ll get: [H + ] = 1.995x10 -4, which rounds to 2.0x10 -4.

44. 20.2 Measuring pH Indicators: HIn(aq) ↔ H + + In - (acid) (base) Indicator changes color when goes from left to right.

45. Universal indicator Indicator solutions and pH meters are two ways to measure the pH of a solution.

46. 20.3 Acid-base theories Goals for this section: Compare and contrast acids and bases as defined by 1. 1.Arrhenius 2. 2.Br Ø nstad-Lowry 3. 3.Lewis

47. Arrhenius In 1887 Swedish chemist Svante Arrhenius proposed a way of defining acids and bases. He said that acids are hydrogen-containing compounds that ionize to yield hydrogen ions [H + ], and bases are compounds that ionize to yield hydroxide ions [OH - ]. Acids that contain one ionizable hydrogen, such as nitric acid HNO 3, are called monoprotic. Acids that contain two ionizable hydrogens, such as sulfuric acid, H 2 SO 4, are called diprotic. But not all hydrogens in an acid are ionizable. Only the hydrogens in very polar bonds are ionizable: δ+ δ- H 2 O H-Cl H + (aq) + Cl - (aq)

48. Arrhenius (continued) δ+ δ- H 2 O H-Cl H + (aq) + Cl - (aq) The two red dots for the H on the right both came from the oxygen (hence it’s a coordinate bond). The electron originally belonging to the H has been stolen by the Chlorine (forming an ion).

49. Arrhenius (continued) In contrast, other acids have H’s that are not involved in polar bonds and therefore not easily “donated” to become an H 3 O + molecule. Here’s an example, ethanoic acid. The 3 H’s attached to the left carbon are not in a polar bond, and are not easily ionized. Only the H attached to the O at the right is in a polar bond and is ionizable. This is a carboxylic acid (COOH group) which is a weak acid.

50. Arrhenius Bases Since potassium and sodium hydroxide are highly soluble in water, they form concentrated basic solutions (which are slippery and taste bitter). They are caustic to the skin. Calcium and magnesium hydroxide are only slightly soluble in water so they form dilute basic solutions only. Some common bases FormulaSolubility in water Potassium hydroxideKOHHigh Sodium hydroxideNaOHHigh Calcium hydroxideCa(OH) 2 Low Magnesium hydroxideMg(OH) 2 Low

51. Bronsted-Lowry Acids and Bases The Arrhenius definition of acid/base is not very comprehensive because it doesn’t deal with chemicals that don’t have H + or OH - but nevertheless are acidic or basic, like aqueous solutions of NH 3 or sodium carbonate Na 2 CO 3, which are basic. In 1923 Bronstad (Danish) and Lowry (English) proposed a new definition. – –An acid is a hydrogen-ion donor – –A base is a hydrogen-ion acceptor – –All Arrhenius acids and bases already meet this definition, but additional substances do also: NH 3 (aq) + H 2 O(l) ↔ NH 4 + (aq) + OH - (aq) Ammonia water ammonium hydroxide ion B-L base B-L acid ion makes solution basic

52. Bronsted-Lowry acids and bases (continued) NH 3 (aq) + H 2 O(l) ↔ NH 4 + (aq) + OH - (aq) base acid conjugate conjugate acid base When ammonia dissolves, NH 4 + is the conjugate acid of the base NH 3. A conjugate acid is the particle formed when a base gains an H + ion. A conjugate base is the particle that remains when an acid has donated a H + ion.

53. Bronsted-Lowry Quiz #1 Another example: HCl(g) + H 2 O(l) ↔ H 3 O + (aq) + Cl - (aq) Which one becomes the conjugate acid? H 3 O + (aq) Which one becomes the conjugate base? Cl - (aq) Therefore which one started out as the acid (H+ ion donor)? HCl(g) Which one started out as the base (H+ ion acceptor)? H 2 O(l)

54. Bronsted-Lowry Quiz #2 Try this one H 2 SO 4 (l) + H 2 O(l) ↔ H 3 O + (aq) + HSO 4 - (aq) Which one is the conjugate acid? H 3 O + (aq) Which one is the conjugate base? HSO 4 - (aq) Which one is the original acid (H ion donor)? H 2 SO 4 (l) Which one is the original base (H ion acceptor)? H 2 O(l)

55. Lewis Acids and Bases A third theory for acids and bases was proposed by Gilbert Lewis. His definition focused on the donation or acceptance of a pair of electrons during the reaction. This concept is more general than the Arrhenius or Bronsted-Lowry definitions (they are valid subsets of the Lewis definition). A Lewis acid is a substance that can accept a pair of electrons to form a covalent bond. A Lewis base is a substance that can donate a pair of electrons to form a covalent bond.

56. Lewis acids and bases In this reaction, a hydroxide ion (OH - ) is a Lewis base because it donated a pair of electrons to the H on the left. It is also a Bronsted-Lowry and/or Arrhenius base. The hydrogen ion is a Lewis acid because it accepted a pair of electrons from the hydroxide ion. (And it is also a Bronsted-Lowry and/or Arrhenius acid).

57. Sample problem 20-7 Identify the Lewis acid and the Lewis base in this reaction Ammonia is donating a pair of electrons, so that means it’s the Lewis base (electron pair donator). BF 3 is accepting the pair of electrons, so it’s the Lewis acid here.

58. Summary of three Acid-Base Definitions Acid-Base Definitions TypeAcidBase ArrheniusH + producerOH - producer Bronsted-LowryH + donorH + acceptor LewisElectron-pair acceptor Electron-pair donor

59. 20.4 Strong and Weak Acids and Bases Officially we’re not doing this section, but it does come up in practice, so just a quick introduction. A strong acid is completely ionized in aqueous solution. (Ex: HCl 100% ionized) A weak acid is only slightly ionized in aqueous solution. (Ex: Ethanoic acid < 1% ionized) A strong base is completely ionized in aqueous solution. (Ex: KOH 100% ionized) A weak base is only slightly ionized in aqueous solution. (Ex: NH 3 < 1% ionized)

60. Put this table on your study buddy and know it for the test !!

61. 20.4 Strong and weak acids/bases Just to avoid possible confusion: The words “concentrated” and “dilute” indicate how much of an acid or base is dissolved in solution. These terms refer to the number of moles of the acid or base in a given volume. The words “strong” or “weak” refer to the extent of ionization of an acid or base. They indicate how many of the dissolved particles actually ionize or dissociate into ions. For example, gastric stomach acid is a DILUTE solution of HCl. That means there are a relatively small number of HCl molecules in a given volume. But it is a STRONG acid because all of those molecules have been ionized.

62. Revisit: Acids Neutralize Bases HCl + NaOH  NaCl + H 2 O Neutralization reactions always produce a salt and water. H 2 SO 4 + 2NaOH  Na 2 SO 4 + 2H 2 O 2HNO 3 + Mg(OH) 2  Mg(NO 3 ) 2 + 2H 2 O In general: HX + YOH → XY + H 2 O acid + base salt + water

63. Practice: Neutralization Reactions Titration known concentration of “titrant” Unknown concentration (you are doing a titration to find out this concentration)

66. Example 3: how many mL of 0.45M HCl must be added to 25.0 mL of 1.00 M KOH to neutralize it? 1. 1.Write and balance equation: HCl + KOH → KCl + H 2 O 2. 2.Find the moles of KOH 1.00 mol KOH/Liter x 0.025 Liter = 0.025 mol KOH 3. 3.Now find moles of HCl to neutralize that 0.025 mol KOH x (1 mol HCl/1 mol KOH) = 0.025 mol HCl 4. 4.Now need to find mL: Since Molarity = moles/liter, let’s solve for liters Liters = moles/molarity = 0.0250 moles HCL / 0.45 M HCl = 0.0556 L or 55.6 mL