Properties and Nomenclature of Acids and Bases

1. Chapter 20 Acids and Bases Ch 20

2. Section 20.1 Describing Acids and Bases • Objectives: • List the properties of acids and bases • Name an acid or base when given the formula

3. Properties of Acids • Acid – a compound that produces hydrogen ions (H+) when dissolved in water. • Taste sour or tart. • Aqueous solutions form electrolytes • Change indicators (turns litmus from blue to red). • React with base to form water and a salt. • Examples? Ch 20

4. Properties of Bases • Base – a compound that produces hydroxide ions (OH-) when dissolved in water. • React with acids to form water and a salt. • Taste bitter. • Feel slippery. • Aqueous solutions form electrolytes • Change indicators (turns litmus from red to blue). • Examples?

5. Names and Formulas of Acids • There are two types of acids: • Binary Acid – Hydrogen is chemically combined with an anion. • Oxyacid – Hydrogen is chemically combined with a polyatomic ion. • The general formula for an acid is HX, where X is either a monatomic or polyatomic anion.

6. Binary Acids • When the anion is monatomic, the acid will have a prefix of hydro- and a suffix of -ic. • HF • HBr • H2S • H3N • HCl Ch 20

7. Oxyacids • When the anion is polyatomic, the acid will not have a prefix. The suffix depends on the ending: -ate becomes -ic and -ite becomes -ous. • HNO3 • HNO2 • H2SO4 • H3PO4 • H3PO3

8. Forming Acids • When forming acids, the charges must equal (like when forming ionic compounds). • Hydrobromic acid • Hydroselenic acid • Sulfurous acid • Sulfuric acid • Hydrosulfuric acid

9. Exceptions • There are a couple of exceptions to the acid nomenclature rules. • Hydrogen and cyanide (CN-): This will follow the binary acid rules HCN = hydrocyanic acid. • All other polyatomic ion without oxygen will following this exception. Sudangrass, sorghum, and sorghum-sudangrass hybrids are among a group of plants that produce cyanide, which can poison livestock under certain conditions.

10. Names and Formulas of Bases • Ionic compounds that are bases are named in the same way as any other ionic compound – the name of the cation followed by the name of the anion. • NaOH • Ca(OH)2 • Al(OH)3

11. Section 20.1 Describing Acids and Bases • Did We Meet Our Objectives • List the properties of acids and bases • Name an acid or base when given the formula

12. Section 20.2 Hydrogen Ions and Acidity • Objectives: • Given the hydrogen-ion or hydroxide-ion concentration, classify a solution as neutral, acidic, or basic • Convert hydrogen-ion concentrations into values of pH and hydroxide-ion concentrations into values of pOH.

13. Hydrogen Ions from Water • Water molecules are highly polar and are in continuous motion. Occasionally, the collisions between water molecules are energetic enough to transfer a hydrogen ion from one water molecule to another. • This produces a positively charged hydronium ion and a negatively charged hydroxide ion. Ch 20

14. The reaction in which two water molecules produce ions is called the self-ionizationof water. This reaction can also be written as a simple dissociation. H2O  H+ + OH- • In water, hydrogen ions (H+) are always joined to water molecules as hydronium ions (H3O+).

15. Hydrogen, Hydronium, and Hydroxide Concentrations • [H+] = Hydrogen Concentration (in molarity) • [H3O+] = Hydronium Concentration (in M) • [OH-] = Hydroxide Concentration (in M)

16. The self-ionization of purewater will produce equal concentrations of hydrogen ions [H+] and hydroxide ions [OH-]. When these concentrations are equal, the solution is said to be neutral. [H+] = [OH-] • If the hydrogen concentration, [H+], is larger than the hydroxide concentration, [OH-], the solution is acidic. [H+] > [OH-] • If the hydroxide concentration, [OH-], is larger than the hydrogen concentration, [H+], the solutions is then basic (alkaline). [H+] < [OH-]

17. The product of the concentrations of the hydrogen ions and hydroxide ions in water is called the ion-product constant for water (Kw). Kw = 1.0 x 10-14 M2 [H+] x [OH-] = 1.0 x 10-14M2 Remember, if [H+] = [OH-], then the solution is neutral. When this occurs, both concentrations will each equal 1.0 x 10-7M. [H+] x [OH-] = Kw [H+] = [OH-] = 1.0 x 10-7

18. Ex #1. If the [H+] in a solution is 1.0 x 10-5M, is the solution acidic, basic, or neutral? What is the [OH-] of this solution? [H+] > 1.0 x 10-7M, acidic [H+] x [OH-] = 1.0 x 10-14 M2 [1.0 x 10-5M] x [OH-] = 1.0 x 10-14 M2 [OH-] = 1.0 x 10-9 M [H+] > [OH-], again this is why it is acidic

19. Ex #2. If the hydroxide-ion concentration of an aqueous solution is 1.0 x 10-3M, what is the hydrogen-ion concentration of the solution? Is the solution acidic, basic, or neutral? [H+] x [OH-] = 1.0 x 10-14 M2 [H+] x [1.0 x 10-3M] = 1.0 x 10-14 M2 [H+] = 1.0 x 10-11M [H+] < [OH-], basic Ch 20

20. The pH Concept • Danish scientist Søren Sørensen proposed the widely used pH scale in 1909. The scale ranges from 0 to 14, with neutral solutions having a pH of 7. Ch 20

21. The pH of the solution is the negative logarithm of the hydrogen-ion concentration. pH = -log[H+]

22. Ex #3. What is the pH of a solution with a hydrogen-ion concentration of 1.0 x 10-10M? pH = -log[H+] pH = -log[1.0 x 10-10M] pH = 10

23. Ex #4. What is the pH of a solution with a hydroxide-ion concentration of 0.0010 M? pOH = -log[OH+] pOH = -log[0.0010 M] pOH = 3 pH + pOH = 14 pH = 11 Ch 20

24. In a definition similar to that of pH, the pOH of a solution equals the negative logarithm of the hydroxide-ion concentration. pOH = -log[OH-] • A simple relationship between pH and pOH makes it easy to find either one when the other is known. pH + pOH = 14

25. Ex #5. What is the pH and the pOH of a solution if [OH-] is 4.0 x 10-11M? pOH = -log[OH-] pOH = -log[4.0 x 10-11M] pOH =10.40 pH + pOH = 14 pH + 10.40 = 14 pH = 3.60 Ch 20

26. How do we find the concentration if we known the pH? • When the pH is not a whole number, you will need a calculator with a yx function key to calculate the ion concentration. (y needs to be base 10) • Or, you can use 2nd log feature (antilog). log [H+] = -pH 2nd log (-pH)

27. Ex #6. What is the molarity of [H+] in each solution? Don’t forget the NEGATIVE! • pH = 7.30 • 5.0 x 10-8 M • pH = 1.80 • 1.6 x 10-2 M • pH = 7.05 • 8.9 x 10-8 M • pH = 6.70 • 2.0 x 10-7 M Ch 20

28. Measuring pH • People need to be able to measure the pH of solutions they use: • Swimming pool, soil conditions, medical diagnoses. • Indicators or a pH meter are often used to measure pH.

29. Acid-Base Indicators • An indicator (In) is an acid or a base that undergoes dissociation in a known pH range. • An indicator is a valuable tool because its acid form and base form have different colors in solution. The following illustrates the dissociation of an indicator (In). OH- H+ Acid Form Base Form

30. pH meters • A pH meter is used to make rapid, accurate pH measurements. • A glass electrode and a reference electrode are connected to a millivoltmeter.

31. Section 20.2 Hydrogen Ions and Acidity • Did We Meet Our Objectives? • Given the hydrogen-ion or hydroxide-ion concentration, classify a solution as neutral, acidic, or basic • Convert hydrogen-ion concentrations into values of pH and hydroxide-ion concentrations into values of pOH.

32. Section 20.3 Acid-Base Theories • Objectives: • Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis • Identify conjugate acid-base pairs in acid-base reactions.

33. Arrhenius Acids and Bases • Svante Arrhenius proposed a revolutionary way of defining and thinking about acids and bases. • He said that acids are hydrogen containing compounds that ionize to yield hydrogen ions in aqueous solution. • He also said that bases are compounds that ionize to yield hydroxide ions in aqueous solution.

34. The Arrhenius concept is limited because it applies only to aqueous solutions and allows for only one kind of base-the hydroxide ion. Arrhenius Acids: HA(aq)↔ H+(aq) + A-(aq) Arrhenius Bases: BOH(aq)↔ B+(aq) + OH-(aq) Ch 20

36. Monoprotic acids – acids that contain one ionizable hydrogen. (HCl) HCl + H2O  H3O+ + Cl- • Diprotic acids – contain two ionizable hydrogens. (H2SO4) H2SO4 + H2O  H3O+ + HSO4- HSO4- + H2O  H3O+ + SO42- • Triprotic acids – contain three ionizable hydrogens. (H3PO4) H3PO4 + H2O  H3O+ + H2PO4- H2PO4- + H2O  H3O+ + HPO42- HPO42- + H2O  H3O+ + PO43- Ch 20

37. Brønsted-Lowry Acids and Bases • Johannes Brønsted and Thomas Lowry independently proposed a new definition for acids and bases. • Brønsted-Lowry Acid – hydrogen-ion donor. • Brønsted-Lowry Base – hydrogen-ion acceptor. NH3 + H2O  NH4+ + OH- Base Acid Conj Conj Acid Base Ch 20

38. Ch 20

39. A conjugate acid is the particle formed when a base gains a hydrogen ion. • A conjugate base is the particle that remains when an acid has donated a hydrogen ion. • A conjugate acid-base pair consists of two substances related by the loss or gain of a single hydrogen ion. Ch 20

40. Ch 20

41. Ex #7. Write the equation for the ionization of HNO2 in water and identify the acid/base pairs. HNO2 + H2O  NO2- + H3O+ Acid Base Conj Conj Base Acid Ch 20

42. Lewis Acids and Bases • A third theory of acids and bases was proposed by Gilbert Lewis. Lewis focused on the donation or acceptance of a pair of electrons during a reaction. • 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. Ch 20

43. Ex #8. Identify the Lewis acid and Lewis base in the following: AlCl3 + Cl- AlCl4-. Ch 20

44. Section 20.3 Recap Ch 20

45. Ch 20

46. Section 20.3 Acid-Base Theories • Did We Meet Our Objectives? • Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis • Identify conjugate acid-base pairs in acid-base reactions. 46

47. Section 20.4 Strengths of Acids and Bases • Objectives: • Define strong acids and weak acids • Calculate an acid dissociation constant (Ka) from concentration and pH measurements • Arrange acids by strength according to their acid dissociation constants (Ka) • Arrange bases by strength according to their base dissociation constants (Kb)

48. Strong and Weak Acids and Bases • Acids are classified as strong or weak depending on the degree to which they ionize in water. • Strong acids are completely ionized in aqueous solutions. i.e. HCl, H2SO4, HNO3 • Weak acids ionize only slightly in aqueous solutions. i.e. CH3COOH, HClO

49. An acid dissociation constant (Ka) is the ratio of the concentration of the dissociated (or ionized) form of an acid to the concentration of the undissociated (non-ionized) form. Ka= Ionized Non-ionized • The acid dissociation constant reflects the fraction of an acid in the ionized form. This is why dissociation constants are sometimes called ionization constants.

50. Weak acids have small Ka values. Large Ka values mean more dissociation, thus stronger acid. • Diprotic and triprotic acids lose their hydrogens one at a time. Each ionization reaction has a separate dissociation constant.