1. Chapter 14 Acids and Bases

2. Homework • Assigned Questions and Problems (odd only) • Section 14.1 (optional) • Section 14.2 Read this section • Section 14.3 Read this section • Section 14.4 33, 35, 37, 39 • Section 14.5 Section on Neutralization Reactions, 43 • Section 14.6 51, 53, 55, 97, 103 • Section 14.8 65, 67, 69, 109, 119 • Section 14.9 71, 73, 75, 77, 79, 81, 83, 85, 87, 107, 121

3. 14.2 Acids: Properties and Examples • Among the most common and important compounds known • Aqueous solutions are important in • biological systems • chemical industrial processes • Early known characteristics • Causes a sour taste (lemons and vinegar) • Causes litmus (dye) to change from blue to red • Dissolves some metals generating hydrogen gas (e.g., zinc and iron) Acetic AcidHC2H3O2(aq) 2Al(s)+ 6HCl(aq)→ 2AlCl3(aq)+3H2(g)

4. 14.2 Acids: Properties and Examples • Common Acids (common laboratory reagents) • Sulfuric acid (H2SO4): most produced chemical in U.S. • Used to produce phosphoric acid and phosphate fertilizers, rust removal in iron and steel production, paper production • Hydrochloric acid (HCl): steel production, organic compound production, pH control and neutralization, main component of stomach acid • Nitric Acid (HNO3): mainly used in the production of fertilzers, explosives, and used to dissolve metals • See table 14.1

5. Section 5.9 Naming Binary Acids • Use the prefix hydro- before the root name of the element • Add the suffix-ic and the word acid to the root name for the element • Example: HCl • hydrochloricacid • Example: HI • hydroiodicacid

6. Section 5.9 Naming Oxyacids • Produce a hydrogen ion (H+) and a polyatomic ion when dissolved in water • Composed of hydrogen, oxygen, and another nonmetal • Use the root name of the polyatomic ion • If it ends in -ate use the suffix -ic acid • If it ends in -ite use the suffix -ous acid • Example: H2SO4 (from SO42- , sulfate ion) • sulfuric acid • Example: H2SO3 (from SO32- , sulfite ion) • sulfurous acid

7. 14.3 Bases: Properties and Examples • Early known characteristics • Causes a bitter taste (when dissolved in water) • Causes litmus (dye) to change from red to blue • Dissolves fats that are placed in base solutions • Feel slippery like soap • Also recognized as an important group of compounds • Aqueous solutions are important in • biological systems • chemical industrial processes

8. Naming Bases • Most bases are usually ionic compounds • The hydroxide ion has a charge of (-1) and is combined with a positively charged ion (group IA or IIA metal ion) • Hydroxides (bases) are named by their cation first, then the word “hydroxide” is added to the metal cation • Most common bases: • NaOH (sodium hydroxide) • KOH (potassium hydroxide) • Ca(OH)2 (calcium hydroxide) • NH4OH (ammoniumhydroxide)

9. 14.4 The Arrhenius Definition • Arrhenius (1884) first person to recognize the essential nature of acids and bases • Defined them in terms of chemical composition • Arrhenius defined an acid as a substance that produces hydrogen ions (H+) in water • Acids are covalent compounds that ionize in water to produce H+ • Acids: hydrogen chloride gas

10. 14.4 Molecular Definitions of Acids and Bases • Arrhenius Acids • When dissolved in water will generate H+ and an anion • The H+ that is generated will give a sour taste • Vinegar (acetic acid) • Lemon (citric acid) • Most acids are oxyacids (will also gen. H+)

11. Types of Acids • Multiprotic Acids • Can donate more than one proton • H2SO4 • H3PO4 Strong Acid Weak Acid

12. Types of Acids • Binary Acids • Molecular compounds in which hydrogen is attached to a second nonmetallic element • Formed from the pure compound which has been dissolved in water • Oxyacids • Molecular compounds composed of hydrogen, oxygen, and a nonmetal (e.g. S, C, N, Cl, or P) • Molecular compounds which become acids when dissolved in water • Acid hydrogen is attached to an oxygen • Formulas are like polyatomic ions with hydrogen(s) attached

13. Types of Acids O O H H C C C C 3 3 O- O H H+ • Organic Acids • Organic compounds (carbon-containing compounds) with acidic properties. • Most common type are called carboxylic acids. The acidic portion of the molecule is called a carboxyl group • For example, acetic acid (the active component of vinegar) contains a carboxyl group Acetate Anion Acetic Acid

14. 14.3 Bases: Properties and Examples • Arrhenius Bases • Arrhenius defined a base as a substance which contains hydroxide and produces hydroxide ions (OH-) in water • Bases are ionic compounds that produce hydroxide ions (OH-) in water • Bases:

15. 14.3 Bases: Properties and Examples • Arrhenius Bases • When dissolved in water will generate OH- and a metal ion. Two common examples: • Bases are sometimes called alkalis • The OH- that is generated will give a bitter taste • Slippery feel (like soap) • Most bases that are generated are composed from group 1A and 2A metals

16. 14.4 The Brønsted-Lowry Definition of Acids and Bases • Arrhenius definition is widely used, but model has limitations • Only for aqueous solutions • Does not explain how compounds such as ammonia (does not contain OH-) produce a basic water solution • New model (1923) by Brønsted and Lowry proposed a broader definition • New model applied to both aqueous and non-aqueous solutions • Explained how compounds without OH- could produce a basic solution when added to water

17. 14.4 The Brønsted-Lowry Definition of Acids and Bases • Brønsted-Lowry Acid • Any substance that can donate a proton (H+) to another substance: a proton donor • Brønsted-LowryBase • Any substance that can accept a proton (H+) from another substance: a proton acceptor • Proton donation does not occur unless a proton acceptor is present • Arrhenius Acids/Bases vs. Br.-Lowry Acids/Bases • All acids in the Arrhenius definition are also acids according to the Bronsted/Lowry model • The Br.-Lowry definition of bases is quite different from that of Arrhenius • H+ ionsproduced by Br.-Lowry acids also react with water

18. Water as a Base • For example: HCl is a Bronsted/Lowry acid because it produces H+ in water • H+ does not exist in water due to the strong attraction to the polar water molecule • In an aqueous solution the H+ ion generated reacts with water to form H3O+ H+ acceptor Hydronium ion H+ donor

19. 14.4 Identifying Brønsted-Lowry Bases and Their Conjugates • Any chemical reaction that involves a Bronsted-Lowry acid must also involve a Br.-Lowry base • The reaction involves a proton transfer • Acid can lose its proton to form a conjugatebase (something that could accept a proton back again) • Base accepts a proton to form a conjugate acid (something that could donate a proton) Conjugate Acid Acid Base Conjugate Base

20. 14.4 Identifying Brønsted-Lowry Bases and Their Conjugates • HCl is the Brønsted-Lowry acid because it is donating a proton (H+) to the water molecule • The water molecule is the Brønsted-Lowry base since it accepts a proton (H+) Base Conjugate Base Conjugate Acid Acid

21. 14.4 Conjugate Acid/Base Pairs • Acids/bases and conj acids/conj bases are related to each other by donating/accepting a single proton (H+) • The acid of the pair has the proton • The base of the pair does not • Every acid has a conjugatebase • Every base has a conjugate acid • “Conjugate” is given to the part of the pair that is produced in the reaction

22. 14.4 Conjugate Acid/Base Pairs • Each acid is related to a base on the opposite side • On the reactants side (left side) substances are called “acid” or “base” • On the products side (right side) substances are called “conjugate acid” or “conjugate base” Conjugate Acid Conjugate Base Acid Base

23. 14.4 Conjugate Acid/Base Pairs • Which of the following represent conjugate acid-base pairs? • H2O, H3O+ • OH-, HNO3 • H2SO4, SO42- • HC2H3O2, C2H3O2-

24. 14.4 Conjugate Acid/Base Pairs H2O, H3O+ OH-, HNO3 H2SO4, SO42- HC2H3O2, C2H3O2-

25. 14.5 Reactions of Acids and BasesNeutralization Reactions • When Arrhenius acids and bases are mixed, they will react with each other • The acidic properties and the basic properties are destroyed and this means both substances are neutralized • A neutralization is the reaction between an acid and a hydroxide base which forms a salt and water

26. 14.5 Reactions of Acids and BasesNeutralization Reactions • A neutralization is a double-displacement reaction • Whenever an acid is completely reacted with a base, a neutralization occurs • The net ionic equation is water formation Acid Base Salt Water

27. 14.5 Reactions of Acids and BasesNeutralization Reactions • Neutralization: The reaction between an acid and a base to form a salt and water • H+ from the acid combines with the OH- from the base to form water • The properties of both reactants are neutralized • The salt contains the positive ion from the base and the negative ion from the acid

28. 14.5 Reactions of Acids and BasesNeutralization Reactions • The reaction hydrochloric acid and sodium hydroxide: • A double replacement reaction • Ionic Equation • Net Ionic Equation water base salt acid

29. 14.8 Water: Acid and Base in One (Ionization of Water) • Substances that can act as acids or bases are termed amphoteric • Water can act as an acid or a base • Amphoteric substance H+ acceptor Hydronium ion Base Acid Conj. Acid Conj. Base H+ donor

30. 14.8 Water: Acid and Base in One (Ionization of Water) • Experiments show that in a sample of pure water, a very small percentage has dissociated to produce ions • It involves a proton transfer (a Br/L acid-base rxn) • The net result is an equal amount of H3O+ and OH- Equal amounts produced Base Acid Conj Acid Conj Base

31. 14.8 Water: Acid and Base in One (Ionization of Water) • The acid-base reaction of water with itself is called autoionization (self-ionization) • It is an equilibrium reaction • Ionization of water to form hydronium ion and hydroxide ion is balanced by the recombination of ions to form water • At 25 ºC, the concentration of each ion: 1.0×10-7 M • Square brackets around a compound denotes “concentration” in moles per liter

32. 14.8 Ion-Product Constant for Water • At any given temperature, the product of the concentration of H+ and OH- is always a constant • This value can be calculated at 25 °C since the concentration (each ion) is known • Kw (Ion-Product Constant for Water) is the product of the H3O+ and OH- ion molar concentrations in water • Valid in pure water or water with solutes

33. 14.8 Ion-Product Constant for Water • If the [H3O+] is increased by the addition of acidic solute, the [OH-] must decrease until the expression is 1.0 × 10-14 is satisfied • Or, if [OH-] is increased by the addition of a basic solute to the water, the [H3O]+ must similarly decrease

34. 14.8 Ion-Product Constant for Water • In an aqueous solution, neither [H3O+] or [OH-] is ever zero • An acid is a substances that will increase the [H3O+] in solution • All acidic solutions have a higher [H3O+] than [OH-] • An base is a substance that will increase the [OH-]ions in solution • All basic solutions have a higher [OH-] than [H3O+]

35. 14.8 Ion-Product Constant for Water • Neutral Solution • Acidic Solution • Basic Solution • In all cases:

36. Using Kw in Calculations, Example 1 • Calculate [H3O+] in a solution in which [OH-] = 2.0x10-2 M. Is this solution acidic, basic or neutral?

37. 14.9 The pH and pOH Scale • H3O+ (from H+) concentrations range from very high values to extremely small valves • Difficult and inconvenient to work with numbers over such a large range • For example, [H3O+]of10 M is 1000 trillion times greater than 10-14 M • The pH scale of a solution was proposed as an easier and more practical way to handle such large numbers

38. 14.9 The pH Scale • The pH scale is defined as the negative logarithm of the molar hydronium ion concentration • Logarithms are exponents • The negative powers of 10 in the concentrations are converted to positive numbers

39. 14.9 The pH Scale • The pH scale uses a common logarithm based on powers of 10 • Expressed mathematically: • For a number expressed in scientific notation with a coefficient of 1, the logarithm of that number is equal to the value of the exponent • Take the log of the number and then multiply by negative one (change the sign)

40. 14.9 Calculating the pH from [H3O+] • The negative log of the H3O+ concentration • To determine the number of significant figures for logs: The number of decimal places for the log is equal to the number of sig. figs. in the concentration (original number) Given: 2 SF 2 D.P.

41. 14.9 Calculating the pH from [H3O+] • Calculate the pH for the following solutions • A solution in which [H3O+]=1.0x10-3 M • A solution in which [OH-]= 5.0x10-5 M

42. Calculating the pH from [H3O+]Calculating the pH from [OH-]

43. Calculating the [H3O+] from pH • It is often necessary to calculate the hydronium ion concentration for a solution from its pH value • The logarithm must be undone • To do this requires determining the antilog of the pH value • The antilog can be obtained on a calculator using the antilog function • Use (-) pH value inv log

44. Calculating the [H3O+] from pH • The pH of rainwater in a polluted area was found to be 3.50. What is the [H3O+] for this rainwater?

45. The pH Scale • pH is a log scale • A change in one unit on the pH scale means a tenfold increase or decrease in [H3O+] • Every time an exponent changes by one, the pH changes by one • Lowering the pH increases the [H3O+] • low pH value= acidic solution • high pH value= basic solution

46. Measuring pH • Use a pH meter • Electronic device that measures the pH of the solution • Use pH paper • Paper has a chemical in it that changes to different colors depending on the pH of the solution • Indicators • A compound that exhibits different colors depending on the pH of its surroundings

47. The pOH Scale • The negative log is also a way of expressing the magnitudes of other quantities • The concentration of OH- can be expressed as pOH • Same as pH scale, but associated with the [OH-] • Low pOH value means high [OH-] • High pOH value means low [OH-]

48. pH and pOH • In an aqueous solution, the sum of the pH and pOH is always 14 • The value 14 corresponds to the of Kw equation -log 1.0 × 10-14

49. pH and pOH [H+] > [OH-] [H+] = [OH-] [H+] < [OH-]

50. Calculating pOH from pH • A sample of rain in an area with severe air pollution has a pH of 3.5. What is the pOH of this rain water?