Acids and Bases

1. Acids and Bases

2. Properties of Acids • An acid is any substance that releases hydrogen ions, H+, into water. • Blue litmus paper turns red in the presence of hydrogen ions. Blue litmus is used to test for acids. • Acids have a sour taste; lemons, limes, and vinegar are acidic.

3. Properties of Bases • A base is a substance that releases hydroxide ions, OH –, into water. • Red litmus paper turns blue in the presence of hydroxide ions. Red litmus is used to test for bases. • Bases have a slippery, soapy feel. • Bases also have a bitter taste; milk of magnesia is a base.

4. Arrhenius Acids and Bases • Svante Arrhenius proposed the following definitions for acids and bases in 1884: • An Arrhenius acid is a substance that ionizes in water to produce hydrogen ions. • An Arrhenius base is a substance that ionizes in water to release hydroxide ions. • For example, HCl is an Arrhenius acid and NaOH is an Arrhenius base.

5. Arrhenius Acids in Solution • All Arrhenius acids have a hydrogen atom bonded to the rest of the molecule by a polar bond. This bond is broken when the acid ionizes. • Polar water molecules help ionize the acid by pulling the hydrogen atom away: HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq) (~100%) HC2H3O2(aq) + H2O(l) → H3O+(aq) + C2H3O2–(aq) (~1%) • The hydronium ion, H3O+, is formed when the aqueous hydrogen ion attaches to a water molecule.

6. Arrhenius Bases in Solution • When we dissolve Arrhenius bases in solution, they dissociate giving a cation and a hydroxide anion. • Strong bases dissociate almost fully and weak bases dissociate very little: NaOH(aq) → Na+(aq) + OH–(aq) (~100%) NH4OH(aq) → NH4+(aq) + OH–(aq) (~1%)

7. Brønsted-Lowry Acids & Bases • The Brønsted-Lowry definitions of acids and bases are broader than the Arrhenius definitions. • A Brønsted-Lowry acid is a substance that donates a hydrogen ion to any other substance. It is a proton donor. • A Brønsted-Lowry base is a substance that accepts a hydrogen ion. It is a proton acceptor.

8. Brønsted-Lowry Acids & Bases • Lets look at two acid/base reactions: • HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) • HCl(aq) + NH3(aq) → NH4Cl(aq) • HCl donates a proton in both reactions and is a Brønsted-Lowry acid. • In the first reaction, the NaOH accepts a proton and is the Brønsted-Lowry base. • In the second reaction, NH3 accepts a proton and is the Brønsted-Lowry base.

9. Strong & Weak Electrolytes • An aqueous solution that is a good conductor of electricity is a strong electrolyte. • An aqueous solution that is a poor conductor of electricity is a weak electrolyte. • The greater the degree of ionization or dissociation, the greater the conductivity of the solution.

10. Electrolyte Strength • Weak acids and bases and insoluble ionic compounds are weak electrolytes. • Strong acids and bases and soluble ionic compounds are strong electrolytes.

12. Strengths of Acids • Acids have varying strengths. • The strength of an Arrhenius acid is measured by the degree of ionization in solution. • Ionization is the process where polar compounds separate into cations and anions in solution. • The acid HCl ionizes into H+ and Cl– ions in solution.

13. Strengths of Bases • Bases also have varying strengths. • The strength of an Arrhenius base is measured by the degree of dissociation in solution. • Dissociation is the process where cations and anions in an ionic compound separate in solution. • A formula unit of NaOH dissociates into Na+ and OH– ions in solution.

14. Strong & Weak Arrhenius Acids • Strong acids ionize extensively to release hydrogen ions into solution. • HCl is a strong acid and ionizes nearly 100% • Weak acids only ionize slightly in solution. • HF is a weak acid and ionizes only about 1%

15. Strong & Weak Arrhenius Bases • Strong bases dissociate extensively to release hydroxide ions into solution. • NaOH is a strong base and dissociates nearly 100% • Weak bases only ionize slightly in solution. • NH4OH is a weak base and only partially dissociates

16. Amphiprotic Compounds • A substance that is capable of both donating and accepting a proton is an amphiprotic compound. • NaHCO3 is an example: • HCl(aq) + NaHCO3(aq) → NaCl(aq) + H2CO3(aq) • NaOH(aq) + NaHCO3(aq) → Na2CO3 (aq) + H2O(l) • NaHCO3 accepts a proton from HCl in the first reaction and donates a proton to NaOH in the second reaction.

17. The pH Scale • A pH value expresses the acidity or basicity of a solution. • Most solutions have a pH between 0 and 14. • Acidic solutions have a pH less than 7. • As a solution becomes more acidic, the pH decreases. • Basic solutions have a pH greater than 7. • As a solution becomes more basic, the pH increases.

18. Acid/Base Classifications of Solutions • A solution can be classified according to its pH: • Strongly acidic solutions have a pH less than 2. • Weakly acidic solutions have a pH between 2 and 7. • Weakly basic solutions have a pH between 7 and 12. • Strongly basic solutions have a pH greater than 12. • Neutral solutions have a pH of 7.

19. Neutralization Reactions • Recall, an acid neutralizes a base to produce a salt and water. • HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) • The reaction produces the aqueous salt NaCl. • If we have an acid with two hydrogens (sulfuric acid, H2SO4), we need two hydroxide ions to neutralize it. • H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + 2 H2O(l)

20. Predicting Neutralization Reactions • We can identify the Arrhenius acid and base that react in a neutralization reaction to produce a given salt such as calcium sulfate, CaSO4. • The calcium must be from calcium hydroxide, Ca(OH)2, and the sulfate must be from sulfuric acid, H2SO4. • H2SO4(aq) + Ca(OH)2(aq) → CaSO4(aq) + 2 H2O(l)

21. Acid-Base Indicators • A solution that changes color as the pH changes is an acid-base indicator. • Three common indicators are methyl red, bromothymol blue, and phenolphthalein. • Each has a different color above and below a certain pH. Chapter 15

22. Acid-Base Indicators • Shown below are the three indicators at different pH values Methyl Red Bromothymol Blue Phenolphthalein Chapter 15

23. Acid-Base Titrations • A titration is used to analyze an acid solution using a solution of a base. • A measured volume of base is added to the acid solution. When all of the acid has been neutralized, the pH is 7. One extra drop of base solution after the endpoint increases the pH dramatically. • When the pH increases above 7, phenolphthalein changes from colorless to pink indicating the endpoint of the titration.

24. Titration Problem • Consider the titration of acetic acid with sodium hydroxide. A 10.0 mL sample of acetic acid requires 37.55 mL of 0.223 M NaOH. What is the concentration of the acetic acid? HC2H3O2(aq) + NaOH(aq) → NaC2H3O2(aq) + H2O(l) • We want concentration acetic acid, we have concentration sodium hydroxide. conc NaOH  mol NaOH  mol HC2H3O2 conc HC2H3O2

25. 37.55 mL solution × × = 0.00837 mol HC2H3O2 0.233 mol NaOH 1 mol HC2H3O2 1000 mL solution 0.00837 mol HC2H3O2 1000 mL solution × 1 mol NaOH 10.0 mL solution 1 L solution Titration Problem Continued • The molarity of NaOH can be written as the unit factor 0.233 mol NaOH / 1000 mL solution. = 0.837 M HC2H3O2

26. Another Titration Problem • A 10.0 mL sample of 0.555 M H2SO4 is titrated with 0.233 M NaOH. What volume of NaOH is required for the titration? • We want mL of NaOH, we have 10.0 mL of H2SO4. • Use 0.555 mol H2SO4/1000 mL and 0.233 mol NaOH/1000 mL.

27. 0.555 mol H2SO4 1 mol H2SO4 0.233 mol NaOH 10.0 mL H2SO4 × × × 1000 mL H2SO4 2 mol NaOH 1000 mL NaOH = 49.8 mL NaOH Problem Continued H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + H2O(l) • 49.8 mL of 0.233 M NaOH is required to neutralize 10.0 mL of 0.555 M H2SO4.

28. Acid-Base Standardization • A standard solution is a solution where the concentration is precisely known. • Acid solutions are standardized by neutralizing a weighed quantity of a solid base. • What is the molarity of a hydrochloric acid solution if 25.50 mL are required to neutralize 0.375 g Na2CO3? 2 HCl(aq) + Na2CO3(aq) → 2 NaCl(aq) + H2O(l) + CO2(g)

29. 1 mol Na2CO3 2 mol HCl 0.375 g Na2CO3 × × 105.99 g Na2CO3 1 mol Na2CO3 = 0.00708 mol HCl 0.00708 mol HCl 1000 mL solution 25.50 mL solution × = 0.277 M HCl 1 L solution Standardization Continued

30. Ionization of Water • Water undergoes an autoionization reaction. Two water molecules react to produce a hydronium ion and a hydroxide ion: • H2O(l) + H2O(l) → H3O+(aq) + OH-(aq) or • H2O(l) → H+(aq) + OH-(aq) • Only about 1 in 5 million water molecules is present as ions so water is a weak conductor. • The concentration of hydrogen ions, [H+], in pure water is about 1 × 10-7 mol/L at 25C.

31. Autoionization of Water • Since [H+] is 1 × 10-7 mol/L at 25C, the hydroxide ion concentration, [OH-], must also be 1 × 10-7 mol/L at 25C: • H2O(l) → H+(aq) + OH-(aq) • At 25C • [H+][OH-] = (1 × 10-7)(1 × 10-7) = 1.0 × 10-14 • This is the ionization constant of water, Kw.

32. [H+] and [OH-] Relationship • At 25C, [H+][OH-] = 1.0 × 10-14. So, if we know the [H+], we can calculate [OH-]. • What is the [OH-] if [H+] = 0.1 M? • [H+][OH-] = 1.0 × 10-14 • (0.1)[OH-] = 1.0 × 10-14 • [OH-] = 1.0 × 10-13

33. The pH Concept • Recall that pH is a measure of the acidity of a solution. • A neutral solution has a pH of 7, an acidic solution has a pH less than 7, and a basic solution has a pH greater than 7. • The pH scale uses powers of ten to express the hydrogen ion concentration. • Mathematically: pH = –log[H+] • [H+] is the molar hydrogen ion concentration

34. Calculating pH • What is the pH if the hydrogen ion concentration in a vinegar solution is 0.001 M? • pH = –log[H+] • pH = –log(0.001) • pH = –(–3) = 3 • The pH of the vinegar is 3, so the vinegar is acidic.

35. Calculating [H+] from pH • If we rearrange the pH equation for [H+], we get: [H+] = 10–pH • Milk has a pH of 6. What is the concentration of hydrogen ion in milk? • [H+] = 10–pH = 10–6 = 0.000001 M • [H+] = 1 × 10–6M.

36. Advanced pH Calculations • What is the pH of blood with [H+] = 4.8 × 10–8M? • pH = –log[H+] = –log(4.8 × 10–8) = – (–7.32) • pH = 7.32 • What is the [H+] in orange juice with a pH of 2.75? • [H+] = 10–pH = 10–2.75 = 0. 0018 M • [H+] = 2.75 × 10–3M

37. Conclusions • pH is a measure of the acidity of a solution. The typical range for pH is 0 to 14. • Neutral solutions have a pH of 7. • Below are some properties of acids and bases.

38. Conclusions Continued • An Arrhenius acid is a substance that ionizes in water to produce hydrogen ions. • An Arrhenius base is a substance that ionizes in water to release hydroxide ions. • A Brønsted-Lowry acid is a substance that donates a hydrogen ion to any other substance. It is a proton donor. • A Brønsted-Lowry base is a substance that accepts a hydrogen ion. It is a proton acceptor

39. Conclusions Continued • In an aqueous solution, [H+][OH-] = 1.0 × 10-14. This is the ionization constant of water, Kw. • pH = –log[H+] • [H+] = 10–pH • Strong acids and bases are strong electrolytes. They are completely dissociated in solution. • Weak acids and bases are weak electrolytes. They are slightly dissociated in solution.

40. Consider the titration of acetic acid with sodium hydroxide. A 10.0 mL sample of acetic acid requires 37.55 mL of 0.223 M NaOH. What is the concentration of the acetic acid? • A 10.0 mL sample of 0.555 M H2SO4 is titrated with 0.233 M NaOH. What volume of NaOH is required for the titration? • What is the molarity of a hydrochloric acid solution if 25.50 mL are required to neutralize 0.375 g Na2CO3?