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Chapter 14

Chapter 14. Acids and Bases. 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

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Chapter 14

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  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?

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