1. 1 Acid Bases Chemistry

2. 2 Acid-Base Objectives 1. Identify the properties of acids and bases. 2. Classify solutions as acidic, basic, or neutral. 3. Compare the Arrhenius and Bronsted-Lowery models. 4. Relate the strength of an acid or base to its degree of ionization. 5. Compare the strength of conjugate acid-base pairs. 6. Explain the meaning of pH and pOH. 7. Relate pH and pOH to the ion product constant of water. 8. Explain how neutralization reactions are used in acid- base titrations.

3. 3 Properties of Acids and Bases Acids 1. sour taste 2. change color in indicator 3. react with active metals to release H 2 gas 4. react with bases to produce salt and water 5. some conduct electric current

4. 4 Properties of Acids and Bases Bases 1. bitter taste 2. change color in indicator 3. dilute solutions feel slippery 4. react with acids to produce salt and water 5. conduct electric current

5. 5 Nomenclature Bases 1.Most are ionic compounds that contain a metal cation and the hydroxide ion; named as ionic compounds 2.Ammonia, NH 3 +, is an exception Binary Acids (Review) 1. Begins with prefix hydro- combined with the first element 2. Root of the second element follows 3. Add the suffix –ic to the end of the second element in place of the ending

6. 6 Nomenclature Oxyacids (Review) 1. Name is based on the polyatomic oxyanion -ending is changed to –ic if polyatomic ends in –ate -ending is changed to –ous if polyatomic ends in –ite 2. Add # hydrogens to front of name to cancel out the charge of polyatomic General Notes: 1. any time hydrogen is the first element in a formula of a compound, the substance is an acid. 2. monoprotic acids contain only one acidic hydrogen. 3. polyprotic acids have more than one acidic hydrogen

7. 7 -diprotic acid: can donate 2 H atoms (contains 2 ionizable hydrogens) ♦ionizes in 2 steps -triprotic acid: can donate 3 H atoms (contains 3 ioniazable hydrogens) ♦ionizes in 3 steps

8. Properties of Acids & Bases Practice (p 596 #1ac, p 601 # 1ab) 1.Write a balanced chemical equation for the reactions that occur between each pair of reactants: a. magnesium and nitric acid (some metals react to form hydrogen gas) c. calcium carbonate and hydrobromic acid (metal and hydrogen carbonates react to form carbon dioxide) 3.Write the steps in the complete ionization of the following polyprotic acids: a. H 2 Seb. H 3 AsO 4 8

9. 9 Strength of Acids and Bases Strong Acid 1. ionizes completely to form H +, H 3 O + ions 2. strong electrolyte 3. depends on the polarity of H and the element its bonded to. 4. ex: HCl, HNO 3, H 2 SO 4, HBr HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq)

10. 10 Strength of Acids and Bases Weak Acid 1. does not ionize completely; fewer ions in solution 2. weak electrolyte 3. typically the more H’s, the weaker the acid 4. ex: HF, HCN, CH 3 COOH, H 2 CO 3 CH 3 COOH (aq) + H 2 O (l)  H 3 O + (aq) + CH 3 COO - (aq)

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12. 12 Strength of Acids and Bases Strong Bases 1. ionizes completely to form OH - ions 2. strong electrolyte 3. the higher the concentration of OH -, [OH - ], the stronger the base 4. ex: metallic hydroxides, especially those of active metals NaOH (s)  Na + (aq) + OH - (aq) alkaline: solution formed when a base completely dissociates in water to yield aqueous OH - ions

13. 13 Strength of Acids and Bases Weak Bases 1. does not ionize completely to form OH - ions 2. weak electrolyte 3. unrelated to the concentration of OH -, [OH - ] 4. ex: CH 3 NH 2 (bases of ammonia) CH 3 NH 2 (aq) + H 2 O (l)  CH 3 NH 3 + (aq) + OH - (aq)

14. 14 Common Commercial Acids H 2 SO 4 : most common acid; metallurgy, petroleum refinement, fertilizers HNO 3 : rare acid; rubber, explosives, plastics, dyes, pharmeceuticals H 3 PO 4 : animal feed, manufacturing fertilizers, flavoring HCl: stomach acid, remove impurities from iron and steel, cleaning masonry, maintain acidity in pools CH 3 COOH: plastics, fungicide

15. 15 Properties of Acids and Bases Review 1. List the characteristics of acids. 2. List the characteristics of bases. 3. The strength of acids or bases depends on what? 4. Name the following acids: a. HFb. H 2 SO 3 5. Write formulas for the following acids: a. nitricb. hydrosulfuric 6. Name the following base: Ca(OH) 2 7. Write a formula for the following base: lithium hydroxide.

16. 16 Acid/Base Theory Why do acids and bases act the way they do and have these particular properties? There are several theories to explain this. Arrhenius Svante Arrhenius (in 1883) was the first to come up with a model explaining acids and bases Arrhenius model: states in an aqueous solution an -acid is a substance that contains hydrogen and ionizes to produce hydrogen ions HCl(g)  H + (aq) + Cl - (aq)

17. 17 Acid/Base Theory Arrhenius Arrhenius model: states in an aqueous solution an -base is a substance that contains a hydroxide group and dissociates to produce a hydroxide ion NaOH(s)  Na + (aq) + OH - (aq) This came from the idea that if acids and bases can produce electric current, they must form ions in solution. Shortcoming: could not explain ammonia (NH 3 ), which does not include a hydroxide group

18. 18 A more inclusive model that would include ammonia was proposed: Brønsted-Lowry Acids and Bases Brønsted-Lowry model, an acid is a hydrogen-ion (proton) donor and a base is a hydrogen-ion (proton) acceptor -more inclusive, meaning not all Arrhenius acids and bases are also B-L acids and bases HCl + NH 3  NH 4 + + Cl - -the more H’s in an acid, the more that can be donated

19. 19 When a Brønsted-Lowry acid donates a hydrogen ion, a conjugate base is formed. When a Brønsted-Lowry base accepts a hydrogen ion, a conjugate acid is formed. Two substances related to each other by the donating and accepting of a single hydrogen ion are a conjugate acid- base pair. HCl + NH 3  NH 4 + + Cl - acid base conj. acid conj. base

20. 20 Conjugate Acid-Base Pairs What are the conjugate acid-base pairs in the following? acid base conj. acid conj. base There is a relationship between the strength of conjugate acids and bases. -stronger acid  weaker conjugate base stronger base  weaker conjugate acid -proton transfer reactions favor the production of the weaker acid and weaker base the reactants must be much stronger as an acid and base than the products

21. 21 Water Notice that when a base dissolves in water, CH 3 NH 2 gains a proton from the H 2 O: CH 3 NH 2 (aq) + H 2 O (l)  CH 3 NH 3 + (aq) + OH - (aq) base acid Notice that when an acid dissolves in water, H 2 O gains a proton from the HCl : HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) acid base Water is amphoteric, meaning it can act as either an acid or a base.

22. 22 Strength and Concentration When talking about acids and bases, strength is different that concentration. Concentration depends on the number of molecules of acid or base in a solution. -use the terms dilute & concentrated Strength refers to the degree of dissociation of ions. -use the terms weak and strong You can have: -dilute solution of a strong acid or base -concentrated solution of a weak acid or base. -dilute solution of weak acid or base -concentrated solution of strong acid or base

23. 23 Strength and Equilibrium Notice the arrow only points one way: HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) -since strong acids ionize completely, essentially there is no reverse reaction Notice the arrow points both ways: CH 3 COOH (aq) + H 2 O (l) ↔ H 3 O + (aq) + CH 3 COO - (aq) -since weak acids do not ionize completely, the system is in equilibrium equilbrium: when the forward and reverse reactions occur at equal rates and all reactant and products are present in the equlibrium mixture.

24. Acid Base Practice 2 p 596 # 1d, p 601 # 3c, p 599 # 2 ab 1d. Write a balanced chemical equation for potassium hydrogen carbonate and hydrochloric acid. 3c. Write the steps for the complete ionization of H 2 SO 3. 2. Identify the conjugate acid-base pairs in the following reactions: a. NH 4 + (aq) + OH - (aq) ↔ NH 3 (aq) + H 2 O(l) b. HBr(aq) + H 2 O(l) ↔ H 3 O + (aq) + Br - (aq) 24

25. 25 Acid-Base Theory Review 1. Label the conjugate acid-base pairs in the following reactions. There will be 2 pairs: an acid-conjugate base pair and a base-conjugate acid pair. 2.What is the relationship between the strength of conjugate acids and bases? 3.What is the importance of Arrhenius’ work? Of Bronsted- Lowery’s work? 4.Why is water amphoteric? 5.What is the difference between strength and concentration?

26. 26 pH Pure water has equal concentrations of H + and OH - ions. -can sometimes act as an acid and sometimes as a base Self-ionization of Water Acids and bases are not the only providers of OH- and H3O+; water also provides these ions, through self- ionization. self-ionization of water: two water molecules produce a hydronium ion and a hydroxide ion by transfer of a proton

27. 27 ionization constant of water (K w ): at given temperatures, water ionizes to form specific concentrations of OH - and H 3 O +. -the mathematical constant is calculated by: K w = [OH - ][ H 3 O + ] - K w increases as temp increases but we use the value at room temperature: - K w = 1.0 x 10 -14 M 2 because [OH - ] = 1.0 x 10 -7 M and [H 3 O + ] = 1.0 x 10 -7 M at room temp

28. 28 Neutral, Acidic, and Basic Solutions In water, [OH - ] = [H + ] because water is neutral If: [OH - ] > [H + ], the solution is basic [OH - ] < [HO + ], the solution is acidic pH Scale pH: defined as the negative of the common logarithm of the hydronium ion concentration pH = -log[H + ] -stands for the French words, pouvoir hydrogene, meaning ‘hydrogen power’

29. 29 In a neutral solution: [H + ] =1.0 x 10 -7 M (and [OH - ] = 1.0 x 10 -7 M ) Ex: What is the pH of a neutral solution at 298K? pH = -log [H + ] pH = -log[1.0 x 10 -7 ] = -[log 1 + log 10 -7 ] = -[0 + (-7)] = -[-7] = 7.0 (sig figs = the number given with [H + ])

30. 30 pOH: defined as the negative of the common logarithm of the hydroxide ion concentration pH = -log[OH - ] Ex: What is the pOH of a neutral solution at 298K? pOH = -log(1.0 x 10 -7 ) = -(-7.0) = 7.0 If you add the pH and pOH, pH + pOH = 14 If you find one, you can subtract from 14 to get the other.

31. 31 We now have the familiar pH scale (0-14): If pH = 7, the solution is neutral pH > 7, the solution is basic pH < 7, the solution is acidic Note: The pH of many weak acids and bases must be calculated experimentally because they do not ionize completely to form hydronium and hydroxide ions.

32. pH & pOH Practice p611 #19all, p 612 #20all 19. What is the pH of the following solutions? a. [H + ] = 1.0 x 10 -2 M b. [H + ] = 3.0 x 10 -6 M c. [OH - ] = 8.2 x 10 -6 M 20. What is the pH and pOH of the following solutions? a. [OH - ] = 1.0 x 10 -6 M b. [OH - ] = 6.5 x 10 -4 M c. [H + ] = 3.6 x 10 -9 M d. [H + ] = 0.025 M 32

33. We can also calculate [OH - ] & [H + ] from pH. Ex: What are the [OH - ] & [H + ] in a healthy person’s blood that has a pH of 7.40? Assume 298K. Step 1: pH pH = -log [H + ] (the log is base 10) -pH = log [H + ] (multiply both sides by -1) 10 -pH = [H + ] (exponential form of log) 10 -7.4 = [H + ] 0.00000004 M = [H + ] 4.0 x 10 - 8 M = [H + ] 33

34. Step 2: pOH We 1 st must determine the pH pH + pOH = 14.00, so pOH = 14.00 – pH pOH = 14.00 – 7.40 = 6.60 Then, pOH = -log [OH - ] (the log is base 10) -pOH = log [OH - ] (multiply both sides by -1) 10 -pOH = [OH - ] (exponential form of log) 10 -6.6 = [OH - ] 0.000000251 M = [OH - ] 2.51 x 10 -7 M = [OH - ] 34

35. [H + ] & [OH - ] Practice p614 # 21 all 21. Calculate the [OH - ] & [H + ] in each solution. a. pH = 2.37 b. pH = 11.05 c. pH = 6.50 35

36. 36 Determining pH There are several ways to determine pH: 1. Indicators can be used to determine the approximate pH of a solution. -acid-base indicators: compounds whose colors are sensitive to pH ♦either weak acids or weak bases ♦come in many different colors ♦exact pH range colors also changes ~transition interval: pH range over which an indicator changes colors

37. 37 2. pH paper is made from soaking the paper in several indicators; also called litmus paper -it can turn almost any color and is fairly accurate -natural indicators include: tea, beet juice, carrot juice, blueberry juice, boiling rose petals, red cabbage juice 3. pH meter determines the pH of a solution by measuring the voltage between the two electrodes that are placed in the solution -used when an exact number is needed

38. 38 pH Review 1. What is the purpose of indicators? 2. Match the following. a. ____ pH = 7 i. [H 3 O + ] > [OH - ] b. ____ pH < 7ii. [H 3 O + ] < [OH - ] c. ____ pH > 7 iii. [H 3 O + ] = [OH - ] 3. What is the concentration of hydronium and hydroxide ions in pure water at 25 o C? 4. Describe the meaning of pH. 5. What is the pOH if the pH = 5.2?

39. 39 Neutralization Reactions neutralization reaction: reaction in which an acid and a base react in aqueous solution to produce a salt and water. -salt: ionic compound made up a cation from a base and an anion from an acid -type of double displacement reaction Ex: Mg(OH) 2 (aq) + 2HCl (aq)  MgCl 2 (aq) + 2H 2 O (l) base acidsalt

40. Neutralization Reactions Practice p 617 # 29ab 29. Write balanced chemical equations for the following acid-base reactions: a. nitric acid and cesium hydroxide b. hydrobromic acid and calcium hydroxide 40

41. 41 A neutralization reaction is carried out by doing a titration. -titration: method for determining the concentration of a solution by reacting a known volume of solution with a known concentration of solution. ♦ex: unknown acid concentration of known volume with a known concentration of base They are carried out by: 1. Placing a pH meter in the solution of acid (known amount, unknown concentration) 2. Filling a buret with a base of known concentration. 3. The base is slowly added until an equivalence point is reached, reading the pH after each addition

42. 42 equivalence point: stoichiometric point when the number of moles of H + ions equals the number of moles of OH - ions. - see p 619, Figure 19-17 -not all equivalence points are at a pH of 7 ♦depends on the strength of the acid and base Many times, chemists use a chemical dye (indicator) instead of a pH meter. -the indicator changes color at the equivalence point, or end point -see figure 19-18, p 619 for some indicators

43. 43 Neutralization Review 1. What is meant by the end point of a titration? 2. What is an equivalence point? 3. When does a neutralization reaction occur? 4. What are the products of a neutralization reaction? 5. When would scientists use a pH meter over an indicator? 6. Are all equivalence points at a pH of 7? Why or why not? 7. How would you test if a shampoo is really ‘pH balanced”?