CHAPTER 14, Study notes of Chemistry

Download CHAPTER 14 and more Study notes Chemistry in PDF only on Docsity! ONLINE Chemistry HMHScience.com (c ) © Ro ya lty F re e/ Co rb is Acids and Bases CHAPTER 14 BIG IDEA Acids are substances that donate hydrogen ions in aqueous solutions. Bases are substances that accept hydrogen ions in aqueous solutions. SECTION 1 Properties of Acids and Bases SECTION 2 Acid-Base Theories SECTION 3 Acid-Base Reactions Why It Matters Video HMHScience.com GO ONLINE Acids and Bases ONLINE LABS  Is It an Acid or a Base?  Effects of Acid Rain on Plants  I’ve Got a Secret DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A main ideas Acids are identified by their properties. Some acids are useful in industry. The properties of bases differ from those of acids. Arrhenius acids and bases produce ions in solution. Key Terms binary acid Arrhenius acid strong acid oxyacid Arrhenius base weak acid How many foods can you think of that are sour? Chances are that almost all the foods you thought of, like those in Figure 1.1a, owe their sour taste to an acid. Sour milk contains lactic acid. Vinegar, which can be produced by fermenting juices, contains acetic acid. Phosphoric acid gives a tart flavor to many carbonated beverages. Most fruits contain some kind of acid. Lemons, oranges, grapefruits, and other citrus fruits contain citric acid. Apples contain malic acid, and grape juice contains tartaric acid. Many substances known as bases are commonly found in household products, such as those in Figure 1.1b. Household ammonia is an ammonia-water solution that is useful for all types of general cleaning. Sodium hydroxide, Na OH, known by the common name lye, is present in some commercial cleaners. Milk of magnesia is a suspension in water of magnesium hydroxide, Mg(OH)2, which is not very water-soluble. It is used as an antacid to relieve discomfort caused by excess hydrochloric acid in the stomach. Aluminum hydroxide, Al (OH)3, and sodium hydrogen carbonate, NaHCO3, are also bases commonly found in antacids. Properties of Acids and Bases NH3(aq) NaOH NaHCO3 Al(OH)3 (a) Fruits and fruit juices contain acids such as citric acid and ascorbic acid. Carbonated beverages contain benzoic acid, phosphoric acid, and carbonic acid. (b) Many household cleaners contain bases such as ammonia and sodium hydroxide. Antacids contain bases such as aluminum hydroxide. Benzoic acid, HC7H5O2 Sorbic acid, HC6H7O2 Phosphoric acid, H3PO4 Carbonic acid, H2CO3 Citric acid, H3C6H5O7 Ascorbic acid, H2C6H6O6 Common Acids and Bases Figure 1.1 Acids and Bases 449 SECTION 1 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A mAin ideA Some acids are useful in industry. The properties of acids make them important chemicals both in the laboratory and in industry. Sulfuric acid, nitric acid, phosphoric acid, hydrochloric acid, and acetic acid are all common industrial acids. Sulfuric Acid Sulfuric acid is the most commonly produced industrial chemical in the world. More than 37 million metric tons of it are made each year in the United States alone. Sulfuric acid is used in large quantities in petroleum refining and metallurgy as well as in the manufacture of fertilizer. It is also essential to a vast number of industrial processes, including the produc- tion of metals, paper, paint, dyes, detergents, and many chemical raw materials. Sulfuric acid is the acid used in automobile batteries. Because it attracts water, concentrated sulfuric acid is an effective dehydrating (water-removing) agent. It can be used to remove water from gases with which it does not react. Sugar and certain other organic compounds are also dehydrated by sulfuric acid. Skin contains organic compounds that are attacked by concentrated sulfuric acid, which can cause serious burns. Nitric Acid Pure nitric acid is a volatile, unstable liquid. Dissolving the acid in water makes the acid more stable. Solutions of nitric acid are widely used in industry. Nitric acid also stains proteins yellow. The feather in Figure 1.6 was stained by nitric acid. The acid has a suffocating odor, stains skin, and can cause serious burns. It is used in making explosives, many of which are nitrogen- containing compounds. It is also used to make rubber, plastics, dyes, and pharmaceuticals. Initially, nitric acid solutions are colorless; however, upon standing, they gradually become yellow because of slight decomposition to brown nitrogen dioxide gas. Nitric Acid and Proteins Concentrated nitric acid stains a feather yellow. Figure 1.6 Chapter 14452 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A Phosphoric Acid Phosphorus, along with nitrogen and potassium, is an essential element for plants and animals. The bulk of phosphoric acid produced each year is used directly for manufacturing fertilizers and animal feed. Dilute phosphoric acid has a pleasant but sour taste and is not toxic. It is used as a flavoring agent in beverages and as a cleaning agent for dairy equip- ment. Phosphoric acid is also important in the manufacture of detergents and ceramics. Hydrochloric Acid The stomach produces HCl to aid in digestion. Industrially, hydro chloric acid is important for “pickling” iron and steel. Pickling is the immersion of metals in acid solutions to remove surface impurities. This acid is also used in industry as a general cleaning agent, in food processing, in the activation of oil wells, in the recovery of magnesium from sea water, and in the production of other chemicals. Concentrated solutions of hydrochloric acid, commonly referred to as muriatic acid, can be found in hardware stores. It is used to maintain the correct acidity in swimming pools and to clean masonry. Acetic Acid Pure acetic acid is a clear, colorless, and pungent-smelling liquid known as glacial acetic acid. This name is derived from the fact that pure acetic acid has a freezing point of 17°C. It can form crystals in a cold room. The fermentation of certain plants produces vinegars containing acetic acid. White vinegar contains 4% to 8% acetic acid. Acetic acid is important industrially in synthesizing chemicals used in the manufacture of plastics. It is a raw material in the production of food supplements—for example, lysine, an essential amino acid. Acetic acid is also used as a fungicide. mAin ideA The properties of bases differ from those of acids. How do bases differ from acids? You can answer this question by compar- ing the following properties of bases with those of acids. 1. Aqueous solutions of bases taste bitter. You may have noticed this fact if you have ever gotten soap, a basic substance, in your mouth. Taste should NEVER be used to test a substance to see if it is a base, because many bases are caustic; they attack the skin, causing severe burns. 2. Bases change the color of acid-base indicators. As Figure 1.7 shows, an acid-base indicator changes to a different color in a basic solution than in an acidic solution. 3. Dilute aqueous solutions of bases feel slippery. You encounter this property of aqueous bases whenever you wash with soap. Base Indicator pH paper turns blue in the presence of this solution of sodium hydroxide. Figure 1.7 Acids and Bases 453 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A 4. Bases react with acids to produce salts and water. It could also be said that “neutralization” of the base occurs when these two substances react to produce a salt and water. 5. Bases conduct electric current. Like acids, bases form ions in aqueous solutions and are thus electrolytes. Unlike acids, there are no special rules for naming basic compounds. The names of bases follow the conventions for chemical compounds outlined in the chapter “Chemical Formulas and Chemical Compounds.” HouseHoLd ACids And BAses QUESTION Which of the household substances are acids, and which are bases? PROCEDURE Record all your results in a data table. 1. To make an acid-base indicator, extract juice from red cabbage. First, cut up some red cabbage and place it in a large beaker. Add enough water so that the beaker is half full. Then, bring the mixture to a boil. Let it cool, and then pour off and save the cabbage juice. This solution is an acid-base indicator. 2. Assemble foods, beverages, and cleaning products to be tested. 3. If the substance being tested is a liquid, pour about 5 mL into a small beaker. If it is a solid, place a small amount into a beaker and moisten it with about 5 mL of water. 4. Add a drop or two of the red cabbage juice to the solution being tested, and note the color. The solution will turn red if it is acidic and green if it is basic. DISCUSSION 1. Are the cleaning products acids, bases, or neither? 2. What are acid/base character- istics of foods and beverages? 3. Did you find consumer warning labels on basic or acidic products? Red cabbage can be used to make an acid-base indicator by extracting its anthocyanin pigment. MATERIALS • dishwashing liquid, dishwasher detergent, laundry detergent, laundry stain remover, fabric softener, and bleach • mayonnaise, baking powder, baking soda, white vinegar, cider vinegar, lemon juice, soft drinks, mineral water, and milk • fresh red cabbage • hot plate • beaker, 500 mL or larger • beakers, 50 mL • spatula • tap water • tongs SAFETY Wear safety goggles, gloves, and an apron. Chapter 14454 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A A molecule of acetic acid contains four hydrogen atoms. However, only one of the hydrogen atoms is ionizable. The hydrogen atom in the carboxyl group in acetic acid is the one that is “acidic” and forms the hydronium ion. This acidic hydrogen can be seen in the structural diagram in Figure 1.10. Aqueous Solutions of Bases Most bases are ionic compounds containing metal cations and the hydroxide anion, OH-. Because these bases are ionic, they dissociate when dissolved in water. When a base completely dissociates in water to yield aqueous OH- ions, the solution is considered a strong base. Sodium hydroxide, NaOH, is a common laboratory base. It is water-soluble and dissociates as shown by the equation below. NaOH(s) H2O    ―⟶ Na+(aq) + OH-(aq) As you will remember from learning about the periodic table, Group 1 elements are the alkali metals. This group gets its name from the fact that the hydroxides of Li, Na, K, Rb, and Cs all form alkaline (basic) solutions. Not all bases are ionic compounds. A base commonly used in house- hold cleaners is ammonia, NH3, which is molecular. Ammonia is a base because it produces hydroxide ions when it reacts with water molecules, as shown in the equation below. NH3(aq) + H2O(l) →  ← NH 4 + (aq) + OH-(aq) Strength of Bases As with acids, the strength of a base also depends on the extent to which the base dissociates, or adds hydroxide ions to the solution. For example, potassium hydroxide, KOH, is a strong base because it completely dissociates into its ions in dilute aqueous solutions. KOH(s) H2O    ―⟶ K+(aq) + OH-(aq) Strong bases are strong electrolytes, just as strong acids are strong electro- lytes. Figure 1.11 lists some strong bases. Acetic Acid Acetic acid contains four hydrogen atoms, but only one of them is “acidic” and forms the hydronium ion in solution. CHeCK FOr uNDerSTANDiNg Differentiate What is the difference between the strength and the concentra- tion of an acid or base? Figure 1.10 H – H | C | H – O || C – O – H acidic hydrogen Figure 1.11 Common Aqueous BAses Strong bases Weak bases Ca(OH)2 → Ca2+ + 2OH- NH3 + H2O → ← NH 4 + + OH- Sr(OH)2 → Sr2+ + 2OH- C6H5NH2 + H2O → ← C6H5NH 3 + + OH- Ba(OH)2 → Ba2+ + 2OH- NaOH → Na+ + OH- KOH → K+ + OH- RbOH → Rb+ + OH- CsOH → Cs+ + OH- Acids and Bases 457 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A Bases that are not very soluble do not produce a large number of hydroxide ions when added to water. Some metal hydroxides, such as Cu(OH)2, are not very soluble in water, as seen in Figure 1.12. They cannot produce concentrated alkaline solutions. The alkalinity of aqueous solutions depends on the concentration of OH- ions in solution. It is unrelated to the number of hydroxide ions in the undissolved compound. Now consider ammonia, which is highly soluble but is a weak electro- lyte. The concentration of OH- ions in an ammonia solution is relatively low. Ammonia is therefore a weak base. Many organic compounds that contain nitrogen atoms are also weak bases. For example, codeine, C18H21NO3, a pain reliever and common cough suppressant found in prescription cough medicine, is a weak base. Reviewing Main Ideas 1. a. What are five general properties of aqueous acids? b. Name some common substances that have one or more of these properties. 2. Name the following acids. a. HBrO b. HBrO3 3. Write the chemical formulas for the following common bases. a. lithium hydroxide b. sodium hydroxide 4. a. What are five general properties of aqueous bases? b. Name some common substances that have one or more of these properties. 5. a. Why are strong acids also strong electrolytes? b. Is every strong electrolyte also a strong acid? Critical Thinking 6. RELATING IDEAS A classmate states, “All compounds containing H atoms are acids, and all compounds containing OH groups are bases.” Do you agree? Give examples. Insoluble Hydroxides The hydroxides of most d-block metals are nearly insoluble in water, as is shown by the gelatinous precipitate, copper(II) hydroxide, Cu(OH)2, in the beaker on the right. Figure 1.12 Na+(aq) + OH-(aq) Cu(OH)2(s) Chloride ion, Cl- Water molecule, H2O Copper(II) ion, Cu2+ Sodium ion, Na+ Chloride ion, Cl- Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s) Chapter 14458 SECTION 1 FORMATIVE ASSESSMENT DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A CROSS-DISCIPLINARY CONNECTION Acid Water—A Hidden Menace M any people are unaware of the pH of the tap water in their home until they are confronted with such phenomena as a blue ring materializing around a porcelain sink drain, a water heater suddenly giving out, or tropical fish that keep dying. Each of these events could be due to acidic water. Acidic water can also cause the amount of lead in the water to rise. The possibility of lead poisoning from home water supplies is a concern. Many older homes still have lead pipes in their plumbing, though most modern homes use copper piping. Copper pipe joints, however, are often sealed with lead- containing solder. Highly acidic water can leach out both the lead from the solder joints and copper from the pipes themselves, which turns the sink drain blue. In addition, people who are in the habit of filling their tea kettles and coffee pots in the morning without letting the tap run awhile first could be adding copper and lead ions to their tea or coffee. Lead poisoning is of particular concern in young children. The absorption rate of lead in the intestinal tract of a child is much higher than in that of an adult, and lead poisoning can permanently impair a child’s rapidly growing nervous system. The good news is that lead poisoning and other effects of acidic water in the home can be easily prevented by following these tips: 1. Monitor the pH of your water on a regular basis, especially if you have well water. This can easily be done with pH test kits (see photograph) that are sold in hardware or pet stores—many tropical fish are intolerant of water with a pH that is either too high (basic) or too low (acidic). The pH of municipal water supplies is already regulated, but regularly checking your water’s pH yourself is a good idea. 2. In the morning, let your water tap run for about half a minute before you fill your kettle or drink the water. If the water is acidic, the first flush of water will have the highest concentration of lead and copper ions. 3. Install an alkali-injection pump, a low-cost, low-maintenance solution that can save your plumbing and lessen the risk of lead poisoning from your own water supply. The pump injects a small amount of an alkali (usually potassium carbonate or sodium carbonate) into your water-pressure tank each time your well’s pump starts. This effectively neutralizes the acidity of your water. Question Suppose that you have tested the drinking water in your home and found it to be too acidic. Design a solution to correct this problem. In your solution, explain the steps you will take to make the water less acidic and what evidence you will collect to show that your solution has worked. The pH of your home’s water supply can be easily monitored using a test kit, such as the one shown here. 459 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A A polyprotic acid is an acid that can donate more than one proton per molecule. Sulfuric acid, H2SO4, and phosphoric acid, H3PO4, are examples of polyprotic acids. The ionization of a polyprotic acid occurs in stages. The acid loses its hydrogen ions one at a time. Sulfuric acid ionizes in two stages. In its first ionization, sulfuric acid is a strong acid. It is completely converted to hydrogen sulfate ions, HSO 4 - . H2SO4(l) + H2O(l) → H3O+(aq) + HSO 4 - (aq) The hydrogen sulfate ion is itself a weak acid. It establishes the following equilibrium in solution. HSO 4 - (aq) + H2O(l) →  ← H3O+(aq) + SO 4 2- (aq) All stages of ionization of a polyprotic acid occur in the same solution. Sulfuric acid solutions therefore contain H3O+, HSO 4 - , and SO 4 2- ions. Note that in sulfuric acid solutions, there are many more hydrogen sulfate and hydronium ions than there are sulfate ions. Sulfuric acid is the type of polyprotic acid that can donate two protons per molecule, and it is therefore known as a diprotic acid. Ionizations of a monoprotic acid and a diprotic acid are shown in Figure 2.2. Monoprotic and Diprotic Acids Hydrochloric acid, HCl, is a strong monoprotic acid. A dilute HCl solution contains hydronium ions and chloride ions. Sulfuric acid, H2SO4, is a strong diprotic acid. A dilute H2SO4 solution contains hydrogen sulfate ions from the first ionization, sulfate ions from the second ionization, and hydronium ions from both ionizations. CRITICAL THINKING Explain Using the Brønsted-Lowry definition, is it possible to have an acid without an accompanying base? FIGURE 2.2 HCl + H2O → H3O+ + Cl- H2SO4 + H2O → H3O+ + HSO 4 - HSO 4 - + H2O → ← H3O+ + SO 4 2- Water molecule, H2O Chloride ion, Cl- Hydronium ion, H3O+ Hydronium ion, H3O+ Sulfate ion, SO 4 2- Hydrogen sulfate ion, HSO 4 - Chapter 14462 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A Phosphoric acid is the type of polyprotic acid known as a triprotic acid— an acid able to donate three protons per molecule. The equations for these reactions are shown below. H3PO4(aq) + H2O(l) →  ← H3O+(aq) + H2PO 4 - (aq) H2PO 4 - (aq) + H2O(l) →  ← H3O+(aq) + HPO 4 2- (aq) HPO 4 2- (aq) + H2O(l) →  ← H3O+(aq) + PO 4 3- (aq) A solution of phosphoric acid contains H3O+, H3PO4, H2PO 4 - , HPO 4 2- , and PO 4 3- . As with most polyprotic acids, the concentration of ions formed in the first ionization is the greatest. There are lesser concentrations of the respective ions from each succeeding ionization. Phospho ric acid is a weak acid in each step of its ionization. MAIN IDEA A Lewis acid or base accepts or donates a pair of electrons. The Arrhenius and Brønsted-Lowry definitions describe most acids and bases. Both definitions assume that the acid contains or produces hydro- gen ions. A third acid classification, based on bonding and structure, includes, as acids, substances that do not contain hydrogen at all. This definition of acids was introduced in 1923 by G. N. Lewis, the American chemist whose name was given to electron-dot structures. Lewis’s definition emphasizes the role of electron pairs in acid-base reactions. A Lewis acid is an atom, ion, or molecule that accepts an electron pair to form a covalent bond. The Lewis definition is the broadest of the three acid definitions you have read about so far. It applies to any species that can accept an elec- tron pair to form a covalent bond with another species. A bare proton (hydrogen ion) is a Lewis acid in reactions in which it forms a covalent bond, as shown below. H+(aq) + :NH3(aq) → [H–NH3]+(aq) or [NH4]+(aq) The formula for a Lewis acid need not include hydrogen. Even a silver ion can be a Lewis acid, accepting electron pairs from ammonia to form covalent bonds. Ag+(aq) + 2:NH3(aq) → [H3N–Ag–NH3]+(aq) or [Ag(NH3)2]+ A compound in which the central atom has three valence electrons and forms three covalent bonds can react as a Lewis acid. It does so by accepting a pair of electrons to form a fourth covalent bond, completing an electron octet. Boron trifluoride, for example, is an excellent Lewis acid. It forms a fourth covalent bond with many molecules and ions. Its reaction with a fluoride ion is shown below. : . .   F . . : : . .   F . . :    B    : . .   F . . : + : . .   F . . : - →       : . .   F . . : : . .   F . . :    B    : . .   F . . : : . .   F . . :       - BF3(aq) + F-(aq) → BF 4 - (aq) Acids and Bases 463 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A The Lewis definition of acids can apply to species in any phase. For example, boron trifluoride is a Lewis acid in the gas-phase combination with ammonia. : . .   F . . : : . .   F . . :    B    : . .   F . . : + :   H  . .   N  . .   H  : H → : . .   F . . : : . .   F . . :    B    : . .   F . . : :   H  . .   N  . .   H  : H A Lewis base is an atom, ion, or molecule that donates an electron pair to form a covalent bond. An anion is a Lewis base in a reaction in which it forms a covalent bond by donating an electron pair. In the example of boron trifluoride reacting with the fluoride anion, F- donates an electron pair to boron trifluoride. F- acts as a Lewis base. BF3(aq) + : . .   F . . : – (aq) → BF 4 - (aq) A Lewis acid-base reaction is the formation of one or more covalent bonds between an electron-pair donor and an electron-pair acceptor. Note that although the three acid-base definitions differ, many compounds may be categorized as acids or bases according to all three descriptions. For example, ammonia is an Arrhenius base because OH- ions are created when ammonia is in solution, it is a Brønsted-Lowry base because it accepts a proton in an acid-base reaction, and it is a Lewis base in all reactions in which it donates its lone pair to form a covalent bond. A comparison of the three acid-base definitions is given in Figure 2.3. Reviewing Main Ideas 1. Label each reactant in the reaction below as a proton donor or a proton acceptor and as acidic or basic. H2CO3 + H2O →  ← HCO 3 - + H3O+ 2. For the reaction below, label each reactant as an electron-pair acceptor or electron-pair donor and as a Lewis acid or a Lewis base. AlCl3 + Cl- → AlCl 4 - Critical Thinking 3. ANALYZING INFORMATION For the following three reactions, identify the reactants that are Arrhenius bases, Brønsted-Lowry bases, and/or Lewis bases. State which type(s) of bases each reactant is. Explain your answers. a. NaOH(s) → Na+(aq) + OH-(aq) b. HF(aq) + H2O(l) → F-(aq) + H3O+(aq) c. H+(aq) + NH3(aq) → NH 4 + (aq) FIGURE 2.3 ACID-BASE DEFINITIONS Type Acid Base Arrhenius H+ or H3O+ producer OH- producer Brønsted-Lowry proton (H+) donor proton (H+) acceptor Lewis electron-pair acceptor electron-pair donor Chapter 14464 SECTION 2 FORMATIVE ASSESSMENT DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A The table in Figure 3.1 shows that a very strong acid, such as HClO4, has a very weak conjugate base, ClO 4 - . The strongest base listed in the table, the hydride ion, H-, has the weakest conjugate acid, H2. In aqueous solutions, all of the strong acids are 100% ionized, forming hydronium ions along with their anion. The acids below hydronium ion in Figure 3.1 do not ionize 100% in water. Water will react as an acid if a very strong base, such as hydride ion, is present. Figure 3.1 Relative StRengthS of acidS and BaSeS Conjugate acid Formula Conjugate base Formula hydriodic acid* HI iodide ion I- perchloric acid* HClO4 perchlorate ion ClO 4 - hydrobromic acid* HBr bromide ion Br- hydrochloric acid* HCl chloride ion Cl- sulfuric acid* H2SO4 hydrogen sulfate ion HSO 4 - chloric acid* HClO3 chlorate ion ClO 3 - nitric acid* HNO3 nitrate ion NO 3 - hydronium ion H3O+ water H2O chlorous acid HClO2 chlorite ion ClO 2 - hydrogen sulfate ion HSO 4 - sulfate ion SO 4 2- phosphoric acid H3PO4 dihydrogen phosphate ion H2PO 4 - hydrofluoric acid HF fluoride ion F- acetic acid CH3COOH acetate ion CH3COO- carbonic acid H2CO3 hydrogen carbonate ion HCO 3 - hydrosulfuric acid H2S hydrosulfide ion HS- dihydrogen phosphate ion H2PO 4 - monohydrogen phosphate ion HPO 4 2- hypochlorous acid HClO hypochlorite ion ClO- ammonium ion NH 4 + ammonia NH3 hydrogen carbonate ion HCO 3 - carbonate ion CO 3 2- monohydrogen phosphate ion HPO 4 2- phosphate ion PO 4 3- water H2O hydroxide ion OH- ammonia NH3 amide ion† NH 2 - hydrogen H2 hydride ion† H- Increasing base strength In cr ea si ng a ci d st re ng th * Strong acids † Strong bases Acids and Bases 467 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A Main idea Some substances act as either acids or bases. You have probably noticed that water can be either an acid or a base. Any species that can react as either an acid or a base is described as amphoteric. For example, consider the first ionization of sulfuric acid, in which water acts as a base. H2SO4(aq) + H2O(l) ⎯→ H3O+(aq) + HSO 4 - (aq) acid1 base2 acid2 base1 However, water acts as an acid in the following reaction. NH3(g) + H2O(l) ⎯→  ←⎯ NH 4 + (aq) + OH-(aq) base1 acid2 acid1 base2 Likewise, water is acting as an acid in the reaction shown in Figure 3.2. Thus, water can act as either an acid or a base and is amphoteric. Such a substance acts as either an acid or a base depending on the strength of the acid or base with which it is reacting—if it’s a stronger acid than the substance, the substance will act as a base. If, however, it’s a stronger base than the substance, the substance will act as an acid. Water as an Acid Calcium hydride, CaH2, reacts vigorously with water to produce hydrogen gas. Water acts as an acid in this reaction because the hydride ion is a very strong base. Figure 3.2 O—H Bonds and Acid Strength Each oxyacid of chlorine contains one chlorine atom and one hydrogen atom. They differ in the number of oxygen atoms they contain. The effect of the increasing O—H bond polarity can be seen in the increasing acid strength from hypochlorous acid to perchloric acid. Figure 3.3 H : . . O. . : . . Cl. . : H : . . O. . : . . Cl. . : . . O. . : H : . . O. . : . . Cl : . . O. . : : . . O. . : H : . . O. . : : . . O. . : Cl : . . O. . : : . . O. . : Hypochlorous acid Chlorous acid Chloric acid Perchloric acid Acidity increases CaH2(s ) + 2H2O(l ) ⎯→ Ca(OH)2(aq ) + 2H2(g ) CHeCK FOr uNDerSTANDiNg Identify Identify the conjugate acid and base in the reaction of calcium hydride with water. Chapter 14468 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A C C ClClO OOCl H OOH H H O H N H HH H HBF F F + O O → Chemistry HMHScience.com GO ONLINE –OH in a Molecule Molecular compounds containing –OH groups can be acidic or amphoteric. The covalently bonded –OH group in an acid is referred to as a hydroxyl group. For the compound to be acidic, a water molecule must be able to attract a hydrogen atom from a hydroxyl group. This occurs more easily when the O–H bond is very polar. Any feature of a molecule that increases the polarity of the O–H bond increases the acidity of a molec ular compound. The small, more-electronegative atoms of nonmetals at the upper right in the periodic table form compounds with acidic hydroxyl groups. All oxyacids are molecular electrolytes that contain one or more of these O–H bonds. Such compounds include chloric and perchloric acids. Figure 3.3 (on the previous page) shows the electron-dot formulas of the four oxyacids of chlorine. Notice that all of the oxygen atoms are bonded to the chlorine atom. Each hydrogen atom is bonded to an oxygen atom. Aqueous solutions of these molecules are acids because the O–H bonds are broken as the hydrogen is pulled away by water molecules. The behavior of a compound is affected by the number of oxygen atoms bonded to the atom connected to the –OH group. The larger the number of such oxygen atoms is, the more acidic the compound is. The electronegative oxygen atoms draw electron density away from the O–H bond and make it more polar. For example, chromium forms three different compounds containing –OH groups, as shown below. basic amphoteric acidic Cr(OH)2 Cr(OH)3 H2CrO4 chromium(II) chromium(III) chromic acid hydroxide hydroxide Notice that as the number of oxygen atoms increases, so does the acidity of the compound. Consider also the compounds shown in Figure 3.4. In acetic acid, but not in ethanol, a second oxygen atom is bonded to the carbon atom connected to the –OH group. That explains why acetic acid is acidic but ethanol is not, even though the same elements form each compound. Main idea Neutralization reactions produce water and a salt. Neutralization reactions commonly occur in nature. One example is when an antacid tablet such as milk of magnesia, Mg(OH)2, neutralizes stomach acid, HCl. The products are MgCl2 and H2O. Some neutralization reactions produce gases. Sodium bicarbonate, NaHCO3, and tartaric acid, H2C4H4O6, are two components in baking powder. When water is added, the two compounds produce carbon dioxide. The escaping carbon dioxide causes foods, such as biscuits, to rise. © Je rr y M as on /P ho to R es ea rc he rs , I nc Polarity and Acids Figure 3.4 (a) CH3COOH Acetic acid is acidic. The second oxygen atom on the carbon draws electron density away from the –OH group, making the O–H bond more polar. (b) C2H5OH Ethanol is essentially neutral. It has no second oxygen atom, so the O—H bond is less polar than in acetic acid, and it is a much weaker acid. H : H. . C . . H : H. . C . . H : . . O. . : H Neutralization Reactions Acids and Bases 469 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A Math Tutor Many chemical reactions that occur in water solutions are reactions involving ions. Soluble ionic compounds dissociate into ions when they dissolve, and some molecular compounds, including acids, ionize when they dissolve. An ionic equation represents the species actually present more accurately than an equation that uses full formulas. Writing Equations for Ionic Reactions Problem-Solving TIPS • All soluble ionic compounds are dissociated into ions. Therefore, soluble ionic compounds are shown as the separated ions in the full ionic equation. Strong acids and bases are also shown as the separated ions in the full ionic equation because they are 100% ionized. • Ions that do not take part in the reaction are called spectator ions. In other words, spectator ions stay in solution and will be labeled “(aq)” on both sides of the equation. Eliminating spectator ions reduces the “clutter” of the full ionic equation and produces a net ionic equation that shows only the species that actually react. Sample Problem Write the net ionic equation for the reaction of aqueous ammonium sulfate and aqueous barium nitrate to produce a precipitate of barium sulfate. The balanced formula equation is (NH4)2SO4(aq) + Ba(NO3)2(aq) -→ 2NH4NO3(aq) + BaSO4(s) Rewrite the equation in full ionic form; because ammonium sulfate and barium nitrate are soluble, they are written as separated ions: 2N H 4 + (aq) + S O 4 2- (aq) + Ba2+(aq) + 2N O 3 - (aq) -→ 2N H 4 + (aq) + 2N O 3 - (aq) + BaSO4(s) Eliminating spectator ions, N H 4 + and N O 3 - , yields the net ionic equation: S O 4 2- (aq) + Ba2+(aq) -→ BaSO4(s) Write full and net ionic equations for the reaction that occurs when hydrochloric acid solution is combined with silver nitrate solution. Hydrochloric acid is a strong acid, so it is completely ionized in solution. Silver nitrate is a soluble ionic compound, so its ions are separated in solution. Although most chlorides are soluble, silver chloride is not, so silver chloride will precipitate. The balanced formula equation is HCl(aq) + AgNO3(aq) -→ AgCl(s) + HNO3(aq) The full ionic equation is H3O+(aq) + Cl-(aq) + Ag+(aq) + N O 3 - (aq) -→ H3O+(aq) + N O 3 -(aq) + AgCl(s) Eliminate spectator ions to obtain the net ionic equation: Cl-(aq) + Ag+(aq) -→ AgCl(s) Answers in Appendix E 1. Aqueous copper(II) sulfate reacts with aqueous sodium sulfide to produce a black precipitate of copper(II) sulfide. Write the formula equation, the full ionic equation, and the net ionic equation for this reaction. 2. Write full and net ionic equations for the reaction that occurs when a solution of cadmium chloride, CdCl2, is mixed with a solution of sodium carbonate, Na2CO3. Cadmium carbonate is insoluble. Chapter 14472 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A Summary SECTION 1 Properties of Acids and Bases KEY TERMS • Acids have a sour taste and react with active metals. Acids change the colors of acid-base indicators, react with bases to produce salts and water, and conduct electricity in aqueous solutions. • Bases have a bitter taste, feel slippery to the skin in dilute aqueous solutions, change the colors of acid-base indicators, react with acids to produce salts and water, and conduct electricity in aqueous solution. • An Arrhenius acid contains hydrogen and ionizes in aqueous solution to form hydrogen ions. An Arrhenius base produces hydroxide ions in aqueous solution. • The strength of an Arrhenius acid or base is determined by the extent to which the acid or base ionizes or dissociates in aqueous solutions. binary acid oxyacid Arrhenius acid Arrhenius base strong acid weak acid SECTION 2 Acid-Base Theories KEY TERMS • A Brønsted-Lowry acid is a proton donor. A Brønsted-Lowry base is a proton acceptor. • A Lewis acid is an electron-pair acceptor. A Lewis base is an electron-pair donor. • Acids are described as monoprotic, diprotic, or triprotic depending on whether they can donate one, two, or three protons per molecule, respec- tively, in aqueous solutions. Polyprotic acids include both diprotic and triprotic acids. Brønsted-Lowry acid Brønsted-Lowry base Brønsted-Lowry acid-base reaction monoprotic acid polyprotic acid diprotic acid triprotic acid Lewis acid Lewis base Lewis acid-base reaction SECTION 3 Acid-Base Reactions KEY TERMS • In every Brønsted-Lowry acid-base reaction, there are two conjugate acid-base pairs. • A strong acid has a weak conjugate base; a strong base has a weak conjugate acid. • Proton-transfer reactions favor the production of the weaker acid and weaker base. • The acidic or basic behavior of a molecule containing -OH groups depends on the electronegativity of other atoms in the molecule and on the number of oxygen atoms bonded to the atom that is connected to the -OH group. • A neutralization reaction produces water and an ionic compound called a salt. • Acid rain can create severe ecological problems. conjugate base conjugate acid amphoteric neutralization salt BIG IDEA Acids are substances that donate hydrogen ions in aqueous solutions. Bases are substances that accept hydrogen ions in aqueous solutions. Chapter Summary 473 CHAPTER 14 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A SECTION 1 Properties of Acids and Bases REVIEWING MAIN IDEAS 1. Compare the general properties of acids with the general properties of bases. 2. a. Distinguish between binary acids and oxyacids in terms of their component elements and the systems used in naming them. b. Give three examples of each type of acid. 3. Identify and describe the characteristic properties of five common acids used in industry. Give some examples of the typical uses of each. 4. Although HCl(aq) exhibits properties of an Arrhenius acid, pure HCl gas and HCl dissolved in a nonpolar solvent exhibit none of the properties of an Arrhenius acid. Explain why. 5. a. What distinguishes strong acids from weak acids? b. Give two examples each of strong acids and weak acids. 6. H3PO4, which contains three hydrogen atoms per molecule, is a weak acid, whereas HCl, which con- tains only one hydrogen atom per molecule, is a strong acid. Explain why. 7. a. What compounds are strong Arrhenius bases? b. Give an example of an aqueous solution of a strong base and one of a weak base. PRACTICE PROBLEMS 8. Name each of the following binary acids: a. HCl b. H2S 9. Name each of the following oxyacids: a. HNO3 c. HClO3 b. H2SO3 d. HNO2 10. Write formulas for the following binary acids and common bases: a. hydrofluoric acid c. sodium bicarbonate b. hydriodic acid d. aluminum hydroxide 11. Write formulas for the following oxyacids: a. perbromic acid b. chlorous acid c. phosphoric acid d. hypochlorous acid SECTION 2 Acid-Base Theories REVIEWING MAIN IDEAS 12. Distinguish between a monoprotic, a diprotic, and a triprotic acid. Give an example of each. 13. Which of the three acid definitions is the broadest? Explain. PRACTICE PROBLEMS 14. a. Write the balanced equations that describe the two-step ionization of sulfuric acid in a dilute aqueous solution. b. How do the degrees of ionization in the two steps compare? 15. Dilute HCl(aq) and KOH(aq) are mixed in chemically equivalent quantities. Write the following: a. formula equation for the reaction b. full ionic equation c. net ionic equation 16. Repeat item 15, but mix H3PO4(aq) and NaOH(aq). 17. Write the formula equation and net ionic equation for each of the following reactions: a. Zn(s) + HCl(aq) -→ b. Al(s) + H2SO4(aq) -→ 18. Write the formula equation and net ionic equation for the reaction between Ca(s) and HCl(aq). SECTION 3 Acid-Base Reactions REVIEWING MAIN IDEAS 19. Define and give an equation to illustrate each of the following substances: a. a conjugate base b. a conjugate acid Review Interactive Review HMHScience.com GO ONLINE Review Games Concept Maps Chapter 14474 CHAPTER 14 DO NOT EDIT--Changes must be made through “File info” CorrectionKey=NL-A