In Chapter 2 we described how elements are synthesized in the cores of giant stars and dispersed throughout galaxies when the stars explode as supernovas. The dis- persed elements become the building blocks of other stars and planets. As our solar system formed, the inner planets— Mercury, Venus, Earth, and Mars— were rich in nonvolatile elements such as iron, silicon, magnesium, and aluminum. Earth was also rich in oxygen— in the form of stable compounds with these and other ele- ments. These compounds formed the rocks and minerals of Earth’s crust and pro- vided an early atmosphere. As Earth cooled, atmospheric water vapor condensed into torrential rains that filled up depressions in the crust, forming the first oceans.
Could the substances in lifeless rocks and air or dissolved in primordial seas have combined to become the organic building blocks of life— compounds like simple sugars, amino acids, and the molecules that form DNA? No one knows for sure, but in the early 1950s Nobel Prize winner Harold Urey (1893–1981) and his student Stanley Miller (1930–2007) tested this hypothesis by assembling a mix- ture of gases believed to have been present in Earth’s early (prebiotic) atmosphere.
Urey and Miller subjected that mixture of gases to an electric current, simulating the lightning that would have been prevalent on early Earth (Figure 3.1). The resulting chemical reactions produced several amino acids, the building blocks of proteins. Samples from additional experiments by Miller in an apparatus simulat- ing the composition of gases from a volcanic eruption were reanalyzed in 2007, and more than 20 amino acids were detected.
Amino acids have also been found in meteorites, which have bombarded Earth since it formed. Extensive analyses of the Murchison meteorite, which landed in Australia in 1969, showed that it contains many of the same amino acids synthesized in the Miller–Urey experiments. Regardless of whether important precursors to biological compounds were synthesized throughout the early atmo- sphere, formed in the localized environment around volcanic eruptions, or arrived on Earth in meteors, current thinking favors an early Earth with a primitive atmosphere and an ocean of water that provided an environment conducive to the synthesis of the molecules of life.
These scientific theories about the origin of life from simple inorganic mole- cules and the evidence that supports them are the product of research carried out LO1 Use Avogadro’s number and the
definition of the mole in calculations Sample Exercises 3.1, 3.2, 3.3, 3.4, 3.5, 3.6, 3.7
LO2 Write balanced chemical equations that describe chemical reactions
Sample Exercises 3.8, 3.9, 3.10, 3.11 LO3 Use balanced chemical equations to relate the mass of a reactant consumed to the mass of a product formed
Sample Exercises 3.12, 3.13
LO4 Determine an empirical formula from the percent composition of a substance Sample Exercises 3.14, 3.15, 3.16
LO5 Determine a molecular formula from the empirical formula and molar mass of a substance
Sample Exercise 3.17
LO6 Use data from combustion reactions to determine empirical formulas of substances
Sample Exercises 3.18, 3.19
LO7 Determine the limiting reactant in a chemical reaction
Sample Exercises 3.20, 3.21 LO8 Calculate the theoretical and percent yields in a chemical reaction Sample Exercises 3.22, 3.23
Learning Outcomes
FIGURE 3.1 The apparatus used by Miller (shown) and Urey to simulate the synthesis of amino acids in the atmosphere of early (prebiotic) Earth. Discharges between the tungsten electrodes were meant to provide the sort of energy that might have come from lightning.
3.1 Air, Life, and Molecules 85 in the 20th and 21st centuries. In the early 19th century, when molecular science
was in its infancy, it was believed that organic compounds could only be made by living organisms. In addition, the prevailing understanding held that chemical compounds were different only because they had different elemental composi- tions. A serendipitous experiment by Friedrich Wửhler (1800–1882) in 1828 led to the demise of both ideas and to their replacement with a view of matter that is a cornerstone of modern molecular science.
Wửhler attempted to synthesize the compound ammonium cyanate (NH4NCO; Figure 3.2a,b) by using silver cyanate (AgNCO) and ammonium chloride (NH4Cl). The product of his reaction demonstrated none of the chemical or physical properties of cyanates. After several other attempts at synthesis, he established that the white crystals produced by the reaction were urea (H2NCONH2; Figure 3.2c). At that time, urea was known only as a material isolated from urine, clearly the product of a living being. Wửhler famously reported to his scientific mentor Jửns Jacob Berzelius (1779–1848) that he had succeeded in making urea “without the inter vention of a kidney.” This result and others like it ultimately led to the modern definition of organic compounds as introduced in Chapter 2— compounds of carbon and hydrogen, often including other elements— and to the understanding that organic compounds identical to those found in living systems can be made in the laboratory.
But Wửhler’s result had another very important consequence. Look carefully at the formulas for ammonium cyanate and urea. If we write them simply in a way that illustrates their composition, ammonium cyanate is CH4N2O, and urea is exactly the same: CH4N2O. The two different substances have the same atomic composi- tion. Scientists ultimately reasoned that the arrangement of atoms within the mole- cules of each compound must be different. This idea led to the modern statement that chemical substances must be defined by not only the number and kind of atoms in their molecules but also the arrangement of those atoms. Composition and arrangement both matter. Berzelius called urea and ammonium cyanate isomers (literally, “same units”), which is the term still used to describe two or more differ- ent compounds with the same atomic composition.
In this chapter we begin with a molecular view of the physical processes and chemical reactions that may have occurred on early Earth. We explore how the quantities of substances produced and consumed in chemical reactions are related.
Then we examine the analytical methods that reveal the composition of pure sub- stances. Finally, we consider a class of chemical reactions that consume oxygen and a fuel, liberating energy— including the energy necessary for life. We will concen- trate on the quantitative aspects of identifying chemical substances and their reac- tions, but keep Wửhler’s result in mind: how the atoms within molecules are arranged matters, too. We will address the important issue of how atoms connect to each other and the resulting shapes of molecules in detail in Chapters 8 and 9.
Chemical Reactions and Earth’s Early Atmosphere
The Earth that formed 4.6 billion years ago was a hot, molten sphere that gradually separated into distinct regions based on differences in density and melting point.
The densest elements, notably iron and nickel, sank to the center of the planet. A less dense mantle, rich in compounds containing aluminum, magnesium, silicon, and oxygen, formed around the core. As time passed and Earth cooled, the mantle fractionated further, allowing a solid crust to form from the components of the mantle that were the least dense and had the highest melting points. The core also
(a) Ammonium ion, NH4+
(b) Cyanate ion, CNO–
(c) Urea, H2NCONH2
FIGURE 3.2 (a, b) Ball-and-stick and space-filling models of the ions in ammonium cyanate and (c) a molecule of urea.
isomers compounds with the same molecular formula but different arrangements of the atoms in their molecules.
separated into a solid inner core and a molten outer core. Figure 3.3 shows the ele- mental compositions of these layers.
Earth’s early crust was torn by the impact of asteroids and widespread volcanic activity. The gases released by these impacts and eruptions generated a primitive atmosphere with a chemical composition different from that of the air we breathe now. Current research favors a view of the early atmosphere of Earth being nearly devoid of molecular oxygen (O2) yet rich in oxygen-containing compounds, includ- ing carbon dioxide (CO2), carbon monoxide (CO), and water vapor (H2O). Minor components of the atmosphere were hydrogen (H2), methane (CH4), and hydrogen sulfide (H2S). Oxides of nitrogen and additional CO may have arisen as a result of heating of the atmosphere during bombardment by large meteorites, whereas traces of other volatile oxides, including sulfur dioxide (SO2) and sulfur trioxide (SO3), may have resulted from volcanic activity. Today, the most abundant gases released by volcanoes like Japan’s Mount Ontake are water vapor, CO2, and SO2 (Figure 3.4).
Sometimes these compounds in Earth’s prebiotic atmosphere combined to make substances with more elaborate molecular structures. For example, sulfur trioxide gas and water vapor are the reactants that combine to form liquid sulfu- ric acid, H2SO4, as a product (Figure 3.5). We use the formulas of these sub- stances in a chemical equation to describe the reaction:
SO31g2 1H2O1g2 SH2SO41,2 (3.1)
Reactants Sulfur
trioxide Water Sulfuric
acid Product
+
FIGURE 3.5 In this combination reaction, a molecule of SO3 and a + molecule of H2O form a molecule of H2SO4.
Crust
O
Fe
O Si Mg Fe Ca
Ni
Si Al Fe Ca Na K Mg
49.5%
88.6%
43.7%
21.6%
16.6%
13.2%
2.1%
8.5%
25.7%
7.5% 4.7% 3.4% 2.6% 2.4% 1.9%
Mantle
Outer core
Inner core FIGURE 3.3 Earth is composed of a solid
inner core, consisting mostly of nickel and iron, surrounded by a molten outer core of similar composition. A rocky mantle, composed mostly of oxygen, silicon, magnesium, and iron, lies between the outer core and a relatively thin solid crust.
FIGURE 3.4 Mount Ontake, the second- highest volcano in Japan (after Mount Fuji), erupted violently on September 27, 2014.
The most abundant gas released in this eruption was water vapor.
3. 2 The Mole 87 The reaction between SO3 and H2O is an example of a combination reaction,
a reaction where two (or more) substances combine to form one product. An important feature of any chemical equation is that it is balanced: every atom that is present in the reactants is also present in the products. This conservation of atoms means that there is also a conservation of mass: the sum of the masses of the reactants always equals the sum of the masses of the products.
The sulfuric acid that formed in Earth’s early atmosphere eventually fell to the planet’s surface as highly acidic rain. This rain landed on a crust made up mostly of metal oxides and metalloid oxides, including the mineral hema- tite, Fe2O3. When that sulfuric acid mixed with hematite, another chemical reaction took place— one that produced slightly water-soluble iron(III) sulfate, Fe2(SO4)3, and liquid water. This reaction is described by the following chem- ical equation: