Edvantage Science AP Chemistry 2 WorkbookPLUS Chapter 6.5 by Edvantage Science

6.5

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6.5 Non-metal and Metal Oxides in Water Warm Up 1. Recall some of the information about chemical bonding that you have studied. Consider the electronegativity values of the main-group elements in groups 1 and 2 and 13 through 17, as shown in the chart below.

Electronegativities of the Elements

2.1

Li Be 1.0

1.5

2.0

Na Mg 0.9

C N O F Ne

2.5

3.0

Al Si P

1.2

1.5

1.8

2.1

3.5

4.0

S Cl Ar

2.5

3.0

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

0.8

1.0

1.3

1.5

1.6

1.5

1.8

1.9

1.6

1.8

2.0

2.4

2.8

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 0.8

1.0

1.3

1.4

1.6

2.2

1.8

2.2

1.9

1.7

1.8

1.9

2.1

2.5

Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 0.7

0.9

1.1

1.3

1.5

1.7

1.9

2.2

2.4

1.9

1.8

1.9

2.0

2.2

Fr Ra Ac 0.7

0.9

1.1

Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 1.1

1,1

1.2

1.1

1.2

1.1

1.2

Th Pa U 1.3

1.5

1.7

(a) Elements from which two of these seven groups will form the most ionic oxides? _____________________________________________________________________________________ (b) In which region of the periodic table are elements located that will form the most covalent oxides? _____________________________________________________________________________________ 2. Consider the chemical equations below, representing the typical reaction of two different oxide compounds with water: Na2O(s) + H2O(l) → 2 NaOH(aq)

SO3(g) + H2O(l) → H2SO4(aq) Determine the type of chemical bonds in each of the reactant oxide compounds and complete the following general statements: (a) In general, the reaction of ____________ (ionic or covalent) oxides with water will produce basic solutions. (b) In general, the reaction of ____________ (ionic or covalent) oxides with water will produce acidic solutions.

Basic Anhydrides

The periodic table summarizes an enormous amount of empirical data about the elements by classifying them according to their chemical and physical properties. One of the most important of these classifications is the designation of an element as either a metal or a non-metal. Within each of these categories, the elements and the compounds they form often behave in characteristic and predictable ways in several important types of chemical reactions.

Chapter 6 Applications of Acid-Base Reactions 415

One example of such a chemical change is the reaction between two of the most common compounds on our planet — the oxides of the elements and water. In general, metal oxides react with water to form bases, and non-metal oxides react with water to form acids. As with most general statements in chemistry, there are exceptions. For a metal oxide to react with water and produce a basic solution, it must be both highly ionic and soluble. These criteria are only met by the oxides formed from the group 1 alkali metals and all but one of the group 2 alkaline earth metals. These families are so-named because their oxides (except for Be) react with water to produce alkaline or basic solutions. Metal oxides that react with water in this way are called basic anhydrides (“anhydride” means “without water”). Although the group 2 metal oxides are less soluble than the alkali metal oxides, their aqueous solutions are basic. Metal oxides formed from the group 1 alkali metals and the group 2 alkaline earth metals (except Be) react with water to produce basic solutions. These oxides are referred to as basic anhydrides.

When basic anhydrides react with water, the metal ions are actually spectators. It is really the dissociated oxide ions that react with water. If you locate the oxide ion on the Table of Relative Strengths of Brønsted-Lowry Acids and Bases (Table A5), you will notice that it is found on the bottom right corner of the table where it is identified as a strong base. Because of its enormous affinity for protons, each oxide ion in water will remove a proton from a water molecule. This converts both itself and the water into hydroxide ions. We can demonstrate the initial dissociation of a metal oxide and the subsequent reaction of its oxide ion with water (written as “HOH”) by using sodium oxide as an example: Dissociation: Net ionic equation: Overall reaction:

Na2O(s) → 2 Na+(aq) + O2–(aq) O2–(aq) + HOH(l) → OH−(aq) + OH−(aq) Na2O(s) + H2O(l) → 2 Na+(aq) + 2 OH−(aq)

Not all metal oxides are considered to be basic anhydrides. Metal oxides are solid at room temperature, but many are not soluble in water. The oxide ions are locked so tightly in the crystal lattice structure of these compounds that they cannot react with water to generate hydroxide ions. There are also several metal oxides that are referred to as amphoteric. This means that, depending on the reaction conditions, they can behave as either acidic oxides or basic oxides. Some of these compounds include the group 2 metal oxide, BeO, and also Cr2O3, Al2O3, Ga2O3, SnO2, and PbO2. Finally, some transition metal oxides in which the metal has a high oxidation number actually act as acidic oxides. (You will learn about oxidation numbers in the next chapter.) For example, manganese(VII) oxide and chromium(VI) oxide both react with water to produce acids: Mn2O7(s) + H2O(l) → 2 HMnO4(aq) permanganic acid CrO3(s) + H2O(l) → H2CrO4(aq)

416 Chapter 6 Applications of Acid-Base Reactions

chromic acid

Quick Check 1. What two criteria must be met for a metal oxide to be considered a basic anhydride? _________________________________________________________________________________________ 2. Which two families in the periodic table contain elements that satisfy these criteria? _________________________________________________________________________________________ 3. Calcium oxide, or quicklime, is a component of dry Portland cement, which, when mixed with water, will eventually harden into concrete. Explain why continual contact of your hands with fresh Portland cement can lead to irritation. _________________________________________________________________________________________

Acidic Anhydrides

Non-metal oxides that react with water are known as acidic anhydrides. In their reactions with water, these molecular or covalent oxides produce acids that are also molecular compounds. In such a reaction, because no oxide ions are released into water, no hydroxide ions are produced in the aqueous solution. Instead, the water molecule binds to the molecular oxide, forming a molecular acid. Figure 6.5.1 and Figure 6.5.2 show the Lewis structures of two gaseous non-metal oxides as they react with water to form acids.

S O

H O S O H O

SO3(g)

H2O(l)

H2SO4(aq)

Figure 6.5.1 The formation of sulfuric acid from sulfur trioxide and water.

O N

O H

H O N

O N2O5(g)

H2O(l)

HNO3(aq)

Figure 6.5.2 The formation of two nitric acid molecules from dinitrogen pentaoxide and water

Acidic anhydrides are normally those containing non-metals with relatively high oxidation states such as SO3¸ N2O5, and Cl2O7. For now, we can interpret that to simply mean molecules containing at least two oxygen atoms per non-metal atom. The “lower” non-metal oxides, such as NO and CO, do not react with water to form acids and so are not considered to be acidic anhydrides. Non-metal oxides that react with water to produce acids are called acidic anhydrides.

As we discussed earlier, several metal oxides react with water to produce acidic solutions, but no non-metal oxides that react with water are known to produce basic solutions. Although carbon monoxide does not combine with water to form an acid, carbon dioxide does. The reaction of CO2 and water produces the weak acid carbonic acid, which is too unstable © Edvantage Interactive 2018

Chapter 6 Applications of Acid-Base Reactions 417

to be isolated in its pure form. Therefore, aqueous solutions of CO2 contain varying amounts of the bicarbonate and hydronium ions and can be represented as shown below. This reaction explains why pure water will gradually become acidic when exposed to air containing CO2 and why unpolluted rainwater is slightly acidic with a pH of about 5.6. CO2(g) + 2 H2O(l)



H3O+(aq) + HCO3−(aq)

In addition to making general statements about the reactions of metal oxides and non-metal oxides with water, we can also detect a periodic trend in the acid-base properties of the element oxides of the main group elements. In general, as elements become less metallic, their oxides that react with water produce more acidic solutions. This occurs as we move both left to right across a chemical period and bottom to top up a chemical family.

Quick Check 1. A student adds a few drops of universal indicator to a small beaker of pure water and sees the expected pale green color indicating a pH of slightly less than 7. She then repeatedly blows exhaled air through a straw into the solution and notices the green color slowly change to yellow and then to pale orange. Explain these results. _________________________________________________________________________________________ 2. Why are no hydroxide ions formed when covalent oxides react with water? _________________________________________________________________________________________

Sample Problem 6.5.1 — Metal and Non-metal Oxides Two 500. mL aqueous solutions on a lab bench have had their labels removed. One solution contains 0.50 mol of Li2O and one solution contains 0.50 mol of SO3. How could phenolphthalein be used to identify each solution? Include the chemical equations as part of your answer.

What to Think About

How to Do It

1. One of the solutes is an alkali metal oxide. This solution must be basic.

The Li2O reacts with water to produce a basic solution of LiOH according to the following:

2. One of the solutes is a covalent oxide that reacts with water. This solution must be acidic.

Li2O(s) + H2O(l) → 2 LiOH(aq)

3. Phenolphthalein is pink in a basic solution.

Phenolphthalein will be pink in this solution. The SO3 reacts with water to form a strong acid according to the following: SO3(g) + H2O(l) → H2SO4(aq) Phenolphthalein will be colorless in this solution.

418 Chapter 6 Applications of Acid-Base Reactions

Practice Problems 6.5.1 — Metal and Non-metal Oxides 1. Complete and balance the following formula equations: (a)

K2O(s)

(b)

MgO(s)

H2O(l)

→

H2O(l)

→

2. (a) Each of the following third period oxides reacts with water to produce acids. Arrange these oxides in order from least acidic to most acidic. SO3

Cl2O7

SiO2

P4O10

Al2O3

________ < ________ < ________ < ________ < ________ (b) Two of the above oxides react with water to produce strong acids. Write the balanced equations for those reactions below.

3. Complete the following statements using either “ionic” or “molecular” in the blank spaces. (a) In general, ionic oxides react with water to produce ______________ (ionic or molecular) bases. (b) In general, molecular oxides that react with water produce acids. Acids are ______________ (ionic or molecular) compounds.

The Problem of Acid Rain

Many of the advantages we enjoy and often take for granted in our industrialized and mobile society also come with environmental costs. One of the most serious and widespread is the problem of acid precipitation in the form of rain, snow, fog, and even dry deposits on particles. As noted earlier, normal rainwater has a pH of about 5.6 because of dissolved CO2. If the pH of rainwater is below 5.3, it is referred to as acid rain. Although natural causes also exist, virtually all of the acid precipitation that humankind is responsible for can ultimately be traced back to the burning of fossil fuels in our homes, vehicles, and power plants, and the processing of mineral ores in industry. The products of such activities and the reactions they undergo are further examples of the important impact chemistry has on our lives. Precipitation in any form is an atmospheric event that involves water. Because our atmosphere is

The Human Causes of a gaseous solution, it should come as no surprise that the major substances responsible for acid precipitation are themselves gases that react with water. Each is a non-metal oxide that reacts with Acid Rain water to produce an acid. Let’s discuss how our modern society produces each of the gaseous substances that cause acids to fall from the sky.

Sources of Sulfur Oxides

Many coal and oil deposits contain sulfur impurities. When electrical power plants burn these fossil fuels to drive steam turbines, the sulfur is oxidized to SO2. For example, the main sulfur impurity in coal is iron disulfide or pyrite, FeS2. The combustion of the coal also burns the pyrite and produces SO2 according to the following reaction: 4 FeS2(s) + 11 O2(g) → 2 Fe2O3(s) + 8 SO2(g) Although the levels of sulfur impurities are often less than 3%, the huge quantities of fossil fuels

Chapter 6 Applications of Acid-Base Reactions 419

burned worldwide result in the production of hundreds of millions of tonnes of SO2. Another source of SO2 is the process of smelting, which extracts metals from their ores by heating. Several important metals such as copper, lead, and zinc are present in Earth’s crust, mainly as sulfides. The first step in the smelting of these minerals usually involves roasting, in which the sulfides are heated at high temperatures in air. For example, when ZnS is heated to about 700°C, the following reaction occurs in the roaster: 2 ZnS(s) + 3 O2(g) → 2 ZnO(s) + 2 SO2(g) Much like carbonic acid, the weak acid produced from the reaction of SO2 and water, called sulfurous acid, H2SO3, is unstable and doesn’t exist in its molecular form in water. As a result, we can represent aqueous solutions of sulfur dioxide as follows: SO2(g) + 2 H2O(l)



H3O+(aq) + HSO3−(aq)

This means that rainwater falling through gaseous SO2 will be more acidic than normal when it reaches the ground. The problem becomes worse in polluted air containing ozone, O3, and fine dust particles where, especially in the presence of sunlight, oxygen and ozone will oxidize some of the SO2 to SO3: 2 SO2(g) + O2(g) → 2 SO3(g) SO2(g) + O3(g) → SO3(g) + O2(g) Sulfur trioxide now reacts with water to form the strong acid sulfuric acid: SO3(g) + H2O(l) → H2SO4(aq) Because 100% ionization occurs when sulfuric acid reacts with water, rainwater in the presence of this acid is much more acidic. A gaseous mixture of SO2 and SO3 is sometimes referred to as SOx. Natural processes such as volcanic eruptions also introduce enormous amounts of sulfur oxides into the atmosphere, but they represent only about 5% to 10% of the total released by human activities. The air we breathe is composed of about 79% nitrogen and 21% oxygen. These two gases will not

Sources of Nitrogen react with each other under normal conditions of temperature and pressure because the activation energy for the reaction is far too high. However, when air is mixed with fossil fuels and burned in Oxides the internal combustion engine of vehicles or in electric power plants, the high temperatures and pressures generated cause the two atmospheric gases to react according to the following reaction: N2(g) + O2(g) → 2 NO(g) When the NO gas in a hot exhaust stream encounters cooler outside air, it reacts with oxygen gas to produce NO2 gas. This gaseous mixture of NO and NO2 is often referred to as NOx. Nitrogen dioxide gas reacts with water in the atmosphere to produce the strong acid HNO3 and the weak acid HNO2 as shown below: 2 NO2(g) + H2O(l) → HNO3(aq) + HNO2(aq) As with sulfur oxides, nature also generates a significant amount of nitrogen oxides. Lightning and decaying vegetation are both sources of NO, but we can do nothing to affect those processes. Figure 6.5.3 summarizes the causes and effects of acid rain.

420 Chapter 6 Applications of Acid-Base Reactions

movement of pollutants

SO2

pollutants in cloud water and precipitation

NO2

Figure 6.5.3 Acid rain kills plants, pollutes rivers, lakes, and streams and erodes stone.

Quick Check 1. Identify the three non-metal oxides that react with water in the atmosphere to produce acid precipitation. ________________________________________________________________________________________ 2. How do coal-burning electrical power plants produce the sulfur oxide responsible for acid rain? ________________________________________________________________________________________ 3. Why can the problem of sulfur oxide acid precipitation become worse in polluted air? ________________________________________________________________________________________

Consequences of Acid Rain

Acid precipitation in the form of rain, snow, fog, or dry deposits on particles has been recorded on every continent on Earth — even at the north and south poles! Because the oxides of sulfur and nitrogen that cause acid precipitation are gases that are carried into the atmosphere, prevailing winds can potentially push clouds containing these oxides more than 1500 km from their sources. For example, Scandinavia receives acid precipitation from Germany, England, and countries in eastern Europe. Southern Ontario and Quebec can receive the acid precipitation that originated anywhere in the U.S. industrial belt extending from Chicago to Boston. Measurements taken in Europe and North America confirm how acidic some precipitation can be. Rainwater with a pH of 2.7 (approximately equal to vinegar) has been recorded in Sweden and a pH of 1.8 (almost equal to stomach acid) has been recorded for rain in West Virginia. The effects of acid rain are both serious and widespread. Both aquatic and terrestrial ecosystems can be severely affected when exposed to acidic conditions. Most species of fish die at pH levels below 5, and entire forests have been devastated by acidic precipitation. Obviously, contaminating the water supplies that our species depends on also has serious direct and indirect consequences. The soils in many areas contain aluminum salts that are nearly insoluble in normal groundwater, but begin to dissolve in more acidic solutions. This dissolving soil not only releases the Al3+ ions that are very toxic to fish, but also leaches out many valuable nutrients from the soil, which are lost in the runoff. Figure 6.5.4 shows some examples of acid rain damage.

Chapter 6 Applications of Acid-Base Reactions 421

Many buildings, monuments, and even cemetery headstones contain CaCO3 in the form of either marble or limestone. Carbonate salts dissolve in acids and long-term exposure to acid rain significantly damages such structures. Although not an environmental consequence, the cost of such damage to our society is significant.

Some Reasons for Optimism

Figure 6.5.4 Trees, soil, water,

Figure 6.5.5 This statue

and fish can all be damaged or destroyed by acid rain

shows the damage from years of exposure to acid rain.

Some lakes in regions where acid precipitation occurs are naturally protected because they are bounded by soils rich in limestone. The same reaction that dissolves the calcium carbonate in statues and buildings allows these lakes to resist changes in their pH. As the limestone dissolves, the hydronium ions from acid rain cause more bicarbonate ions to form as shown below. CO32–(aq) + H3O+(aq)



HCO3−(aq) + H2O(l)

Over time, as the [HCO3−] increases, the lakes become effective bicarbonate/carbonate buffer systems and maintain a relatively constant pH. If a lake’s soil does not contain sufficient limestone, a temporary solution is to add CaCO3 or CaO directly to the lake. Over time, however, such lakes will eventually once again suffer the effects of acid rain unless the sources of the SO2 and NO2 are controlled. Growing environmental awareness on a global scale over the past several decades has focused the attention of the public, the politicians who represent them, and the scientific community on attempting to solve the problem of acid precipitation. There is still much work to be done, but there are few disciplines as powerful as chemistry when it comes to solving our problems and so there are reasons for optimism. Strict government emission standards are forcing automobile manufacturers to produce much cleaner and more efficient vehicles. Modern vehicles are equipped with a catalytic converter, which significantly reduces the levels of carbon monoxide, unburned hydrocarbons, and nitrogen oxides released into the atmosphere. One of the reaction chambers in the converter uses catalytic materials such as transition metal oxides and palladium or platinum metals to convert gaseous nitrogen oxides in the exhaust stream to nitrogen and oxygen gas before it leaves the tailpipe. The catalyzed reaction can be represented as shown below. 2 NOx(g) → x O2(g) + N2(g) (The value of “x” is either 1 or 2.) Over the past several years, most major car companies have produced “hybrid” vehicles, which normally use two distinct power sources to move the vehicle. The most common hybrid vehicles are hybrid electric vehicles (HEVs), which combine an internal combustion engine with one or more electric motors. In addition, several major car companies including General Motors and Nissan have

422 Chapter 6 Applications of Acid-Base Reactions

invested heavily in the production of electric vehicles such as the Volt and the Leaf, which are available to the general public. Industrial efforts to reduce sulfur and nitrogen oxide emissions have also intensified over the past several decades. One method of removing sulfur dioxide gas from the exhaust stream of a coal-fired power plant or a metal smelter is by a process called scrubbing. This involves first blowing powdered limestone (CaCO3) into the combustion chamber where heat decomposes it to CaO and CO2. The calcium oxide (lime) then combines with the sulfur dioxide gas to produce solid calcium sulfite: CaO(s) + SO2(g) → CaSO3(s) As a second step, to remove the CaSO3 and any unreacted SO2, an aqueous suspension of CaO is then sprayed into the exhaust gases before they reach the smokestack to produce a thick suspension of CaSO3 called a slurry. The process is not without its drawbacks, however. First, it is expensive and consumes a great deal of energy. Second, because no use has yet been found for the CaSO3, the great quantities of this compound that are produced by the process are usually buried in landfills. A recently developed process for removing SO2 involves using H2S gas to convert the SO2 into elemental sulfur according to the following reaction: 16 H2S(g) + 8 SO2(g) → 3 S8(s) + 16 H2O(l) To reduce the nitrogen oxide emissions from power plants, gaseous ammonia is reacted with the hot stack gases to produce nitrogen and water vapor: 4 NO(g) + 4 NH3(g) + O2(g) → 4 N2(g) + 6 H2O(l) The problem of acid precipitation has not yet been solved. A number of major hurdles must still be overcome, particularly as they relate to international agreements and enforcement. Economics, employment, industrial development, political will, environmental protection, and scientific advancement do not always merge harmoniously on our planet. As you read this, consider the problem a challenge to you and your generation. Although the stakes are very high and the problems are significant, human ingenuity and chemistry have combined to surmount many enormous challenges in our world in the past. Maybe it’s your turn to begin such an endeavor now!

Sample Problem 6.5.2 — Acid Rain Lead is most commonly found in Earth’s crust as lead(II) sulfide, PbS, a mineral called galena. The first step in smelting lead converts the galena to lead(II) oxide by the process of roasting. (a) Write the formula equation for the reaction. (b) Identify the product of the reaction that contributes to acid precipitation.

What to Think About 1. Roasting reacts oxygen in air with the sulfide mineral. 2. The non-metal oxide product is precursor to acid precipitation.

How to Do It

2 PbS(s) + 3 O2(g) → 2 PbO(s) + 2 SO2(g) Both the SO2 and the SO3 that it can produce contribute to acid precipitation.

Chapter 6 Applications of Acid-Base Reactions 423

Practice Problems 6.5.2 — Acid Rain 1. Write the formula equations for the reactions between each of the three oxides and water that result in the production of the acids responsible for acid precipitation.

2. Explain why the release of Al3+ ions into aquatic ecosystems poses a serious threat to the organisms living there

3. How does the process of “scrubbing” remove sulfur dioxide gas from exhaust streams at coal-fired power plants and smelters? Include the relevant chemical equations in your answer.

424 Chapter 6 Applications of Acid-Base Reactions