Introductory Chemistry, 3rd Edition Nivaldo Tro

Introductory Chemistry, 3rd Edition Nivaldo Tro
Chapter 14 Acids and Bases Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA 2009, Prentice Hall

Tro's Introductory Chemistry, Chapter 14
Types of Electrolytes Salts are water-soluble ionic compounds. All strong electrolytes. Acids form H+1 ions in water solution. Bases combine with H+1 ions in water solution. Increases the OH-1 concentration. May either directly release OH-1 or pull H+1 off H2O. Tro's Introductory Chemistry, Chapter 14

Properties of Acids Sour taste. React with “active” metals. I.e., Al, Zn, Fe, but not Cu, Ag or Au. 2 Al + 6 HCl ® 2 AlCl3 + 3 H2 Corrosive. React with carbonates, producing CO2. Marble, baking soda, chalk, limestone. CaCO3 + 2 HCl ® CaCl2 + CO2 + H2O Change color of vegetable dyes. Blue litmus turns red. React with bases to form ionic salts. Tro's Introductory Chemistry, Chapter 14

Common Acids Tro's Introductory Chemistry, Chapter 14

Structure of Acids Oxyacids have acid hydrogens attached to an oxygen atom. H2SO4, HNO3 Tro's Introductory Chemistry, Chapter 14

Properties of Bases Also known as alkalis. Taste bitter. Alkaloids = Plant product that is alkaline. Often poisonous. Solutions feel slippery. Change color of vegetable dyes. Different color than acid. Red litmus turns blue. React with acids to form ionic salts. Neutralization. Tro's Introductory Chemistry, Chapter 14

Common Bases Tro's Introductory Chemistry, Chapter 14

Structure of Bases Most ionic bases contain OH ions. NaOH, Ca(OH)2 Some contain CO32- ions. CaCO3 NaHCO3 Molecular bases contain structures that react with H+. Mostly amine groups. Tro's Introductory Chemistry, Chapter 14

Arrhenius Theory Bases dissociate in water to produce OH- ions and cations. Ionic substances dissociate in water. NaOH(aq) → Na+(aq) + OH–(aq) Acids ionize in water to produce H+ ions and anions. Because molecular acids are not made of ions, they cannot dissociate. They must be pulled apart, or ionized, by the water. HCl(aq) → H+(aq) + Cl–(aq) In formula, ionizable H is written in front. HC2H3O2(aq) → H+(aq) + C2H3O2–(aq) Tro's Introductory Chemistry, Chapter 14

Arrhenius Theory, Continued
HCl ionizes in water, producing H+ and Cl– ions. NaOH dissociates in water, producing Na+ and OH– ions. Tro's Introductory Chemistry, Chapter 14

Brønsted–Lowry Theory
A Brønsted-Lowry acid–base reaction is any reaction in which an H+ is transferred. Does not have to take place in aqueous solution. Broader definition than Arrhenius. Acid is H+ donor; base is H+ acceptor. Since H+ is a proton, acid is a proton donor and base is a proton acceptor. Base structure must contain an atom with an unshared pair of electrons to bond to H+. In the reaction, the acid molecule gives an H+ to the base molecule. H–A + :B  :A– + H–B+ Tro's Introductory Chemistry, Chapter 14

Comparing Arrhenius Theory and Brønsted–Lowry Theory
HCl(aq) + H2O(l)  Cl−(aq) + H3O+(aq) HF(aq) + H2O(l)  F−(aq) + H3O+(aq) NaOH(aq) + H2O(l)  Na+(aq) + OH−(aq) + H2O(l) NH3(aq) + H2O(l)  NH4+(aq) + OH−(aq) Arrhenius theory HCl(aq)  H+(aq) + Cl−(aq) HF(aq)  H+(aq) + F−(aq) NaOH(aq)  Na+(aq) + OH−(aq) NH4OH(aq)  NH4+(aq) + OH−(aq) Tro's Introductory Chemistry, Chapter 14

Amphoteric Substances
Amphoteric substances can act as either an acid or a base. They have both transferable H and an atom with a lone pair. HCl(aq) is acidic because HCl transfers an H+ to H2O, forming H3O+ ions. Water acts as base, accepting H+. HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq) NH3(aq) is basic because NH3 accepts an H+ from H2O, forming OH–(aq). Water acts as acid, donating H+. NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq) Tro's Introductory Chemistry, Chapter 14

Conjugate Pairs In a Brønsted-Lowry acid-base reaction, the original base becomes an acid in the reverse reaction, and the original acid becomes a base in the reverse process. Each reactant and the product it becomes is called a conjugate pair. The original base becomes the conjugate acid; the original acid becomes the conjugate base. Tro's Introductory Chemistry, Chapter 14

Example—Identify the Brønsted–Lowry Acids and Bases and Their Conjugates in the Reaction. H2SO H2O  HSO4– + H3O+ When the H2SO4 becomes HSO4, it loses an H+, so H2SO4 must be the acid and HSO4 its conjugate base. When the H2O becomes H3O+, it accepts an H+, so H2O must be the base and H3O+ its conjugate acid. H2SO H2O  HSO4– + H3O+ Acid Base Conjugate Conjugate base acid Tro's Introductory Chemistry, Chapter 14

Example—Identify the Brønsted-Lowry Acids and Bases and Their Conjugates in the Reaction, Continued. HCO3– H2O  H2CO3 + HO– When the HCO3 becomes H2CO3, it accepts an H+, so HCO3 must be the base and H2CO3 its conjugate acid. When the H2O becomes OH, it donates an H+, so H2O must be the acid and OH its conjugate base. HCO3– H2O  H2CO3 + HO– Base Acid Conjugate Conjugate acid base Tro's Introductory Chemistry, Chapter 14

Practice—Write the Equations for the Following Reacting with Water and Acting as a Monoprotic Acid. Label the Conjugate Acid and Base. HBr HSO4-1 Tro's Introductory Chemistry, Chapter 14

Practice—Write the Equations for the Following Reacting with Water and Acting as a Monoprotic Acid. Label the Conjugate Acid and Base, Continued. HBr + H2O ® Br H3O+1 Acid Base Conjugate Conjugate base acid HBr HSO4-1 HSO H2O ® SO H3O+1 Acid Base Conjugate Conjugate base acid Tro's Introductory Chemistry, Chapter 14

Neutralization Reactions
H+ + OH- H2O Acid + base salt + water Double-displacement reactions. Salt = cation from base + anion from acid. Cation and anion charges stay constant. H2SO4 + Ca(OH)2 → CaSO4 + 2 H2O Some neutralization reactions are gas evolving, where H2CO3 decomposes into CO2 and H2O. H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2 Tro's Introductory Chemistry, Chapter 14

Titration Titration is a technique that uses reaction stoichiometry to determine the concentration of an unknown solution. Titrant (unknown solution) is added from a buret. Indicators are chemicals that are added to help determine when a reaction is complete. The endpoint of the titration occurs when the reaction is complete. Tro's Introductory Chemistry, Chapter 14

Acid–Base Titration The base solution is the titrant in the buret. As the base is added to the acid, the H+ reacts with the OH– to form water. But there is still excess acid present, so the color does not change. At the titration’s endpoint, just enough base has been added to neutralize all the acid. At this point, the indicator changes color: [H+] = [OH–] Tro's Introductory Chemistry, Chapter 14

Example 14.4—What Is the Molarity of an HCl Solution if mL Is required to Titrate mL of M NaOH? NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l) Given: Find: 12.54 mL NaOH, mL HCl M HCl Solution Map: Relationships: M = mol/L, 1 mol NaOH = 1 mol HCl, 1 mL = 0.001L mL HCl L HCl L NaOH mol NaOH M mol HCl mL Solve: Check: The unit is correct, the magnitude is reasonable. Tro's Introductory Chemistry, Chapter 14

Strong Acids The stronger the acid, the more willing it is to donate H. Use water as the standard base. Strong acids donate practically all their Hs. 100% ionized in water. Strong electrolyte. [H3O+] = [strong acid]. [ ] = molarity. HCl ® H+ + Cl- HCl + H2O® H3O+ + Cl- Tro's Introductory Chemistry, Chapter 14

Strong Acids, Continued
Hydrochloric acid HCl Hydrobromic acid HBr Hydroiodic acid HI Nitric acid HNO3 Perchloric acid HClO4 Sulfuric acid H2SO4 Tro's Introductory Chemistry, Chapter 14

Strong Acids, Continued
Pure water HCl solution Tro's Introductory Chemistry, Chapter 14

Weak Acids Weak acids donate a small fraction of their Hs. Most of the weak acid molecules do not donate H to water. Much less than 1% ionized in water. [H3O+] << [weak acid]. HF Û H+ + F- HF + H2O Û H3O+ + F- Tro's Introductory Chemistry, Chapter 14

Weak Acids, Continued Hydrofluoric acid HF Acetic acid HC2H3O2 Formic acid HCHO2 Sulfurous acid H2SO3 Carbonic acid H2CO3 Phosphoric acid H3PO4 Tro's Introductory Chemistry, Chapter 14

Weak Acids, Continued Pure water HF solution Tro's Introductory Chemistry, Chapter 14

Degree of Ionization The extent to which an acid ionizes in water depends in part on the strength of the bond between the acid H+ and anion compared to the strength of the bond between the acid H+ and the O of water. HA(aq) + H2O(l)  A−(aq) + H3O+(aq) Tro's Introductory Chemistry, Chapter 14

Relationship Between Strengths of Acids and Their Conjugate Bases
The stronger an acid is, the weaker the attraction of the ionizable H for the rest of the molecule is. The better the acid is at donating H, the worse its conjugate base will be at accepting an H. Strong acid HCl + H2O → Cl– + H3O+ Weak conjugate base Weak acid HF + H2O  F– + H3O+ Strong conjugate base

Strong Bases The stronger the base, the more willing it is to accept H. Use water as the standard acid. Strong bases, practically all molecules are dissociated into OH– or accept Hs. Strong electrolyte. Multi-OH bases completely dissociated. [HO–] = [strong base] x (# OH). NaOH ® Na+ + OH- Tro's Introductory Chemistry, Chapter 14

Strong Bases, Continued
Lithium hydroxide LiOH Sodium hydroxide NaOH Potassium hydroxide KOH Calcium hydroxide Ca(OH)2 Strontium hydroxide Sr(OH)2 Barium hydroxide Tro's Introductory Chemistry, Chapter 14

Weak Bases In weak bases, only a small fraction of molecules accept Hs. Weak electrolyte. Most of the weak base molecules do not take H from water. Much less than 1% ionization in water. [HO–] << [strong base]. NH3 + H2O Û NH4+ + OH- Tro's Introductory Chemistry, Chapter 14

Autoionization of Water
Water is actually an extremely weak electrolyte. Therefore, there must be a few ions present. About 1 out of every 10 million water molecules form ions through a process called autoionization. H2O Û H+ + OH– H2O + H2O Û H3O+ + OH– All aqueous solutions contain both H3O+ and OH–. The concentration of H3O+ and OH– are equal in water. [H3O+] = [OH–] = 1 x 10-7M at 25 °C in pure water. Tro's Introductory Chemistry, Chapter 14

Ion Product of Water The product of the H3O+ and OH– concentrations is always the same number. The number is called the ion product of water and has the symbol Kw. [H3O+] x [OH–] = 1 x = Kw. As [H3O+] increases, the [OH–] must decrease so the product stays constant. Inversely proportional. Tro's Introductory Chemistry, Chapter 14

Acidic and Basic Solutions
Neutral solutions have equal [H3O+] and [OH–]. [H3O+] = [OH–] = 1 x 10-7 Acidic solutions have a larger [H3O+] than [OH–]. [H3O+] > 1 x 10-7; [OH–] < 1 x 10-7 Basic solutions have a larger [OH–] than [H3O+]. [H3O+] < 1 x 10-7; [OH–] > 1 x 10-7 Tro's Introductory Chemistry, Chapter 14

Ba(OH)2 = Ba2+ + 2 OH– therefore:
Example—Determine the [H3O+] for a M Ba(OH)2 and Determine Whether the Solution Is Acidic, Basic, or Neutral. Ba(OH)2 = Ba OH– therefore: [OH–] = 2 x = = 4.0 x 10−4 M [H3O+] = 2.5 x M. Since [H3O+] < 1 x 10−7, the solution is basic. Tro's Introductory Chemistry, Chapter 14

Practice—Determine the [H3O+] Concentration and Whether the Solution Is Acidic, Basic, or Neutral for the Following: [OH–] = M [OH–] = 3.50 x 10-8 M Ca(OH)2 = 0.20 M Tro's Introductory Chemistry, Chapter 14

Practice—Determine the [H3O+] Concentration and Whether the Solution Is Acidic, Basic, or Neutral for the Following, Continued: [OH–] = M [OH–] = 3.50 x 10-8 M Ca(OH)2 = 0.20 M [H3O+] = 1 x 10-14 2.50 x 10-4 = 4.00 x 10-11 [H3O+] < [OH-1], therefore, base. [H3O+] = 1 x 10-14 3.50 x 10-8 = 2.86 x 10-7 [H3O+] > [OH-1], therefore, acid. [OH-1] = 2 x 0.20 = 0.40 M [H3O+] = 1 x 10-14 4.0 x 10-1 = 2.5 x 10-14 [H3O+] < [OH-1], therefore, base. Tro's Introductory Chemistry, Chapter 14

Complete the Table [H+] vs. [OH-]
Tro's Introductory Chemistry, Chapter 14

pH The acidity/basicity of a solution is often expressed as pH. pH = ─log[H3O+], [H3O+] = 10−pH Exponent on 10 with a positive sign. pHwater = −log[10-7] = 7. Need to know the [H+] concentration to find pH. pH < 7 is acidic; pH > 7 is basic; pH = 7 is neutral. Tro's Introductory Chemistry, Chapter 14

pH, Continued The lower the pH, the more acidic the solution; the higher the pH, the more basic the solution. 1 pH unit corresponds to a factor of 10 difference in acidity. Normal range is 0 to 14. pH 0 is [H+] = 1 M, pH 14 is [OH–] = 1 M. pH can be negative (very acidic) or larger than 14 (very alkaline).

pH of Common Substances
1.0 M HCl 0.0 0.1 M HCl 1.0 Stomach acid 1.0 to 3.0 Lemons 2.2 to 2.4 Soft drinks 2.0 to 4.0 Plums 2.8 to 3.0 Apples 2.9 to 3.3 Cherries 3.2 to 4.0 Unpolluted rainwater 5.6 Human blood 7.3 to 7.4 Egg whites 7.6 to 8.0 Milk of magnesia (saturated Mg(OH)2) 10.5 Household ammonia 10.5 to 11.5 1.0 M NaOH 14

Example—Calculate the pH of a 0
Example—Calculate the pH of a M Ba(OH)2 Solution and Determine if It Is Acidic, Basic, or Neutral. Ba(OH)2 = Ba OH− therefore, [OH-] = 2 x = = 2.0 x 10-3 M. [H3O+] = 1 x 10-14 2.0 x 10-3 = 5.0 x 10-12M pH = −log [H3O+] = −log (5.0 x 10-12) pH = 11.3 pH > 7 therefore, basic. Tro's Introductory Chemistry, Chapter 14

Practice—Calculate the pH of the Following Strong Acid or Base Solutions. M HCl M Ca(OH)2 0.25 M HNO3 Tro's Introductory Chemistry, Chapter 14

Practice—Calculate the pH of the Following Strong Acid or Base Solutions, Continued. M HCl therefore, [H3O+] = M. M Ca(OH)2 therefore, [OH–] = M. 0.25 M HNO3 therefore, [H3O+] = 0.25 M. pH = −log (2.0 x 10-3) = 2.70 [H3O+] = 1 x 10−14 1 x 10−2 = 1.0 x 10−12 pH = −log (1.0 x 10−12) = 12.00 pH = −log (2.5 x 10−1) = 0.60 Tro's Introductory Chemistry, Chapter 14

Example—Calculate the Concentration of [H3O+] for a Solution with pH 3.7. [H3O+] = 10-pH [H3O+] = means < [H+1] < [H3O+] = 2 x 10-4 M = M. Tro's Introductory Chemistry, Chapter 14

Practice—Determine the [H3O+] for Each of the Following:
pH = 2.7 pH = 12 pH = 0.60 Tro's Introductory Chemistry, Chapter 14

Practice—Determine the [H3O+] for Each of the Following, Continued:
pH = 2.7 pH = 12 pH = 0.60 [H3O+] = 10−2.7 = 2 x 10−3 M = M [H3O+] = 10−12 = 1 x M [H3O+] = 10−0.60 = 0.25 M Tro's Introductory Chemistry, Chapter 14

pOH The acidity/basicity of a solution may also be expressed as pOH. pOH = ─log[OH−], [OH−] = 10−pOH Exponent on 10 with a positive sign. pOHwater = −log[10−7] = 7. Need to know the [OH−] concentration to find pOH. pOH < 7 is acidic; pOH > 7 is basic, pOH = 7 is neutral. Tro's Introductory Chemistry, Chapter 14

pOH, Continued The lower the pOH, the more basic the solution; the higher the pOH, the more acidic the solution. 1 pOH unit corresponds to a factor of 10 difference in basicity. Normal range is 0 to 14. pOH 0 is [OH−] = 1 M; pOH 14 is [H3O+] = 1 M. pOH can be negative (very basic) or larger than 14 (very acidic). pH + pOH = Tro's Introductory Chemistry, Chapter 14

Example—Calculate the pH of a 0
Example—Calculate the pH of a M Ba(OH)2 Solution and Determine if It Is Acidic, Basic, or Neutral. Ba(OH)2 = Ba OH− therefore, [OH−] = 2 x = = 2.0 x 10−3 M. pOH = -log [OH−] = -log (2.0 x 10−3) pOH = 2.70 pH = pOH = pH = 11.30 pH > 7 therefore, basic. Tro's Introductory Chemistry, Chapter 14

Practice—Calculate the pOH and pH of the Following Strong Acid or Base Solutions. M KOH M Ca(OH)2 0.25 M HNO3 Tro's Introductory Chemistry, Chapter 14

Practice—Calculate the pOH and pH of the Following Strong Acid or Base Solutions, Continued. M KOH therefore, [OH–] = M. M Ca(OH)2 therefore, [OH–] = M. 0.25 M HNO3 therefore, [H3O+] = 0.25 M. pOH = −log (2.0 x 10-3) = 2.70 pH = – 2.70 = 11.30 pOH = −log (1.0 x 10-2) = 2.00 pH = – 2.00 = 12.00 pH = −log (2.5 x 10-1) = 0.60 pOH = – 0.60 = 13.40 Tro's Introductory Chemistry, Chapter 14

Buffers Buffers are solutions that resist changing pH when small amounts of acid or base are added. They resist changing pH by neutralizing added acid or base. Buffers are made by mixing together a weak acid and its conjugate base. Or weak base and its conjugate acid. Tro's Introductory Chemistry, Chapter 14

How Buffers Work The weak acid present in the buffer mixture can neutralize added base. The conjugate base present in the buffer mixture can neutralize added acid. The net result is little to no change in the solution pH. Tro's Introductory Chemistry, Chapter 14

Acetic Acid/Acetate Buffer

Nonmetal Oxides Are Acidic
Nonmetal oxides react with water to form acids. Causes acid rain. CO2 (g) + H2O(l) → H2CO3(aq) 2 SO2(g) + O2(g) + 2 H2O(l) → 2 H2SO4(aq) 4 NO2(g) + O2(g) + 2 H2O(l) → 4 HNO3(aq) Tro's Introductory Chemistry, Chapter 14

What Is Acid Rain? Natural rain water has a pH of 5.6. Naturally slightly acidic due mainly to CO2. Rain water with a pH lower than 5.6 is called acid rain. Acid rain is linked to damage in ecosystems and structures. Tro's Introductory Chemistry, Chapter 14

What Causes Acid Rain? Many natural and pollutant gases dissolved in the air are nonmetal oxides. CO2, SO2, NO2. Nonmetal oxides are acidic. CO2 + H2O  H2CO3 2 SO2 + O2 + 2 H2O  2 H2SO4 Processes that produce nonmetal oxide gases as waste increase the acidity of the rain. Natural—volcanoes and some bacterial action. Man-made—combustion of fuel. Weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced. Tro's Introductory Chemistry, Chapter 14

Damage from Acid Rain Acids react with metals and materials that contain carbonates. Acid rain damages bridges, cars, and other metallic structures. Acid rain damages buildings and other structures made of limestone or cement. Acidifies lakes affecting aquatic life. Dissolves and leaches more minerals from soil. Making it difficult for trees. Tro's Introductory Chemistry, Chapter 14

Damage from Acid Rain circa 1995 circa 1935 Tro's Introductory Chemistry, Chapter 14

pH of Rain in Different Regions

Sources of SO2 from Utilities

Acid Rain Legislation 1990 Clean Air Act attacks acid rain. Forces utilities to reduce SO2. The result is acid rain in the northeast is stabilized and begins to be reduced. Tro's Introductory Chemistry, Chapter 14

Introductory Chemistry, 3rd Edition Nivaldo Tro

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