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The physical and chemical bases of energy
Chapter 2
The physical and chemical bases of energy 2.1 Energy, work, and power Calories and Joules 2.2 The different forms of energy Chemical energy Radiant energy Heat energy 2.3 The Laws of Thermodynamics The First Law of Thermodynamics Work Enthalpy The Second Law of Thermodynamics Entropy 2.4 Gaia hypothesis 2.5 Carbon and energy The forms of carbon Measures of carbon Carbon chemistry 2.6 Recommended reading
There is no better way to begin the study of ecological energetics than by starting with an understanding of the pertinent definitions and terminology of physics and physical chemistry. Learn this terminology early, become comfortable with the units of measure, know the basic concepts, and bioenergetics will come a lot easier.
2.1 Energy, work, and power Energy is the capacity to do work. The unit of measure for energy is the erg, which is the work performed when a force of one dyne acts through a distance of one centimeter. The unit of force, the dyne, yields to a mass of one gram the acceleration of one centimeter per second (cm s1). Since an erg of energy is such a small quantity, a larger unit, the joule, which is equal to 107 ergs, becomes a more convenient unit of measure. A unit of heat used frequently in The Global Carbon Cycle and Climate Change. https://doi.org/10.1016/B978-0-12-820244-9.00002-0 Copyright © 2020 Elsevier Inc. All rights reserved.
5
6 The Global Carbon Cycle and Climate Change
physical chemistry is the calorie (¼ 4.184 J). The calorie is the heat energy required to raise the temperature of one gram of water from 14.5 oC to 15.5oC. The calorie is defined as being equal to 4.1840 absolute joules. The calorie is a relatively small unit of measure, and for most chemical and biological calculations the kilocalorie (103 calories) is used. The kilocalorie (kcal) is the unit which is typically used in discussing dietary intake and is often written as Calorie. A Calorie equals 103 calories, or a kcal. Calories and Joules. A calorie is the energy needed to raise the temperature of 1 g of water through 1 C (also expressed as 4.1868 J, the unit of energy in the International System of Units). A joule is the energy expended when 1 kg is moved 1 m by a force of 1 Newton (N). Use of joules is now recommended by international convention and is the preferred standard unit to measure heat (FAO, 2003). Nutritionists and food scientists concerned with large amounts of energy generally use kiloJoules (kJ ¼ 103 J) or megaJoules (MJ ¼ 106 J). For many decades, food energy has been expressed in calories, and studies in the field of ecological energetics have traditionally used calories as the measure for energy. In order to retain consistency with research reported in the scientific literature, values used for energy in this book are in calories. The conversion factors for joules and calories are: 1 cal ¼ 4.184 J and 1 J ¼ 0.239 cal.
2.2 The different forms of energy Energy can exist in various forms, but those of greatest importance to living organisms are mechanical, chemical, radiant, and heat energy (Table 2.1). Mechanical energy has two forms: kinetic and potential. Kinetic energy, or free energy, can be described as the “useful energy” which a body possesses by
TABLE 2.1 Units of measure for energy in its various forms and transformations. Energy
Intensity
Mechanical (ergs)
Force (dynes)
Kinetic (ergs) Potential (ergs) Chemical (calories)
Capacity Change in distance (cm) -1
1
Velocity (cm s )
/2 Mass (g)
-2
Height x accereraltion (cm s ) 1
Heat of combustion (cal g ) 2
Mass (g) Mass (g)
Radiant (calories)
Radiation flux (cal cm )
Surface area (cm2)
Heat (calories)
Difference in temperature (o C)
Heat capacity (cal per oC)
Vol. expansion (ergs)
Pressure (dynes cm-2)
Change in volume (cm3)
Electrical (joules)
Difference in potential (volts)
Surface (ergs)
1
Surface tension (ergs cm )
Coulombs (amps x sec.) Change in area (cm2)
The physical and chemical bases of energy Chapter | 2
7
value of its motion, and is measured by the amount of work which is done in bringing that body to rest. Examples would be a moving ball or the Brownian movement of molecules. Potential energy is stored energy, which is only potentially useful until its conversion into the kinetic or free energy where it becomes available to accomplish work. Energy may be stored in a system by virtue of position, as for example, a stone above the Earth’s surface, a steel spring under compression, or by virtue of chemical properties due to the arrangement of atoms and electrons within a molecule. Conversion of energy from the potential form to the kinetic form involves movement, i.e., motion. Chemical energy. All organisms must work to live, and they require a source of potential energy which can be utilized in order to perform the life processes. This energy can be found in the form of the chemical energy of biomass used as food. Energy can also be in the form of the chemical energy of inorganic molecules utilized as an alternative energy source to radiant energy by chemotrophs. Assemblies of atoms in matter can be rearranged into different groups; thus, by the movement of atoms and the creation of different atomic bonds, chemical energy is liberated. The combustion (oxidation) of coal in a furnace or food by the respiratory processes in a cell releases energy which can be used to accomplish work. Both of these processes illustrate the conversion of chemical to mechanical energy. Life processes on this Earth have evolved around carbon chemistry, and most chemical energy sources are derived from organic compounds. However, as we shall see shortly, there are some notable exceptions. Radiant energy. The sun, a vast incandescent sphere of gas, releases energy by the nuclear transmutation of hydrogen to helium, and it is upon this energy source that life on Earth depends. Radiant energy is the energy of electromagnetic radiation. Because electromagnetic radiation can be conceptualized as a stream of photons, radiant energy can be viewed as photon energy. Alternatively, EM radiation can be viewed as an electromagnetic wave, which carries energy in its oscillating electric and magnetic fields. These two views are completely equivalent and are reconciled to one another in quantum field theory. Solar radiation is energy in the form of electromagnetic waves involving a rhythmic exchange between potential and kinetic energy. Electromagnetic radiation can have frequencies, or wave lengths, of different energy content and interactions (e.g., absorptivities) with matter. Heat energy. This is a very special form of energy resulting from the random movements of molecules, which by virtue of their motion, possess kinetic energy. Heat is evolved when all other forms of energy are transformed and work is performed. All work, including the growth and reproduction of living organisms, represent the transformation of energy and ultimately results in the production of heat. For example, when an animal during respiration releases the potential energy of glucose, approximately two-thirds of it is converted
8 The Global Carbon Cycle and Climate Change
into mechanical energy to be used for work (activity and growth) and onethirds is given off as heat. There are instances of work where heat is absorbed (endodermic processes): the cooling unit of a refrigerator or the fixation of atmospheric nitrogen by certain bacteria are examples; but, these processes are not self-supporting energetically. Nitrogen fixation is always accompanied by the exothermic breakdown of organic substrates. Heat energy released by an exothermic process is never used with complete efficiency by the endergonic process, and so whenever work is done the trend is always toward heat production. In natural processes, changes from one form of energy to another (except to heat) are normally incomplete, because the movement, already shown to be necessary for energy conversion, involves either friction or heat production. Temperature is the relative measure which is used to characterize the amount of heat present in a system. Several common systems are used, only two of which concern us: the Celsius and the Kelvin scales. The Celsius (centigrade) scale (oC) establishes zero (0 ) as the freezing point of water and 100 degrees as the boiling point of water. The Kelvin scale is an absolute measure which establishes zero at the temperature (273 C) at which all molecular motion ceases. Because freezing and boiling points limit life processes, we will conventionally utilize the Celsius scale. Heat may be characterized by the properties of two phases: sensible heat and latent heat. The significance of each of these phases will shortly become apparent. Suffice at present to distinguish between the two phases as follows: sensible heat is that which can be measured by an increase in temperature of a body, for example, the warming action of sunlight irradiating a metal plate. Latent heat is the heat absorption by a body without an equivalent increase in temperature, such as the heats of freezing or vaporization of water (heat of vaporization of water ¼ 539 cal at 100 C). Energy flow is expressed as the product of two factors: (1) an intensity factor (or gradient), and (2) a capacity factor (amount). Energy, work, and heat are all expressed in the same units: calories, joules, or ergs. It should be evident that the different energies may be compared, but that no relationship exists between the capacity factors alone. For example, electric energy may be converted into heat energy, but the rise in temperature cannot be calculated from the voltage, unless the number of coulombs and the heat capacity of the system are known. It is also clear that the same quantity of work can be accomplished by a small quantity of water passing through a turbine from a great height as by a large quantity of water passing through from a short distance. Thus, arises one of the fundamental principles of thermodynamicsdthe interconvertibility of energy forms as well as the “trade-offs” between the intensity and capacity factors. All forms of energy are interconvertible; when conversions do occur, they do so according to rigorous laws of exchange. These are the Laws of Thermodynamics.
The physical and chemical bases of energy Chapter | 2
9
2.3 The Laws of Thermodynamics The First Law of Thermodynamics. The First Law of Thermodynamics is also known as the Law of Conservation of Energy, since it defines that the sum of all energies in an isolated system is constant. In other words, energy may be transformed from one state to another, but it can neither be created nor destroyed. The total energy in universe remains constant, but it is continuously becoming more diffuse throughout the universe. Remember that the capacity factor for mechanical and chemical energy is mass (m). Thus, Einstein showed that if there is a change in mass, Dm, Energy ¼ Dmc2,
(2.1)
where: c is the velocity of light (3 1010 cm s1). Therefore, 1 g of water is equivalent to 9 1020 ergs of energy. It should be evident that the Law of Conservation of Energy and the Law of Conservation of Mass are essentially the same, and that no violation of thermodynamics occurs when energy is converted into mass or mass is converted into energy. However, a change in energy of a system will be brought about if the system does work, or if it absorbs or evolves heat. Thus, when a change of any kind occurs in a closed system (where the amount of matter is fixed but energy is able to enter or leave) an increase or decrease occurs in the internal energy (E) of the system itself; heat (q) is evolved or absorbed, and work (w) is done: DE
the decrease in internal energy ðcalÞ
¼
q the heat ðcalÞ given off by the system
w
the work ð-calÞ done by the system
(2.2)
Work. The work performed by a system is the energy transferred by the system to its surroundings. The negative value of work indicates that a positive amount of work done by the system has led to energy being lost from the system. The First Law of Thermodynamics also encompasses the more specific relationship of constant heat sums, which is of considerable importance to biologists interested in energy transformations. It states that the total amount of heat produced, or absorbed, from a chemical reaction which takes place in stages is equal to the total amount of heat evolved, or consumed, when the reaction occurs directly in one step. The evolution of living systems has utilized biochemical mechanisms by which chemical compounds can be reduced in steps, thus enabling more efficient energy capture and utilization to occur. A good biological example is the metabolic oxidation of glucose to carbon dioxide and water: Direct reaction (combustion) C6H12O6 þ 6O2 /6H2O þ 6CO2 þ 673 kcal of energy
10 The Global Carbon Cycle and Climate Change
Two-stage reaction (fermentation) (a) C6H12O6 /2C2H5OH þ 2CO2 þ 18 kcal of energy (b) 2C2H5OH þ 6O2 /6H2O þ 4CO2 þ 655 kcal of energy (a) þ (b) C6H12O6 þ 6O2 /6H2O þ 6CO2 þ 673 kcal of energy Thus, no matter which pathway a particular reaction follows, the total amount of heat evolved, or absorbed, is always the same. There will be more discussion about the biological significance of this phenomenon later. Several other energy relationships are also pertinent to bioenergetics. Enthalpy. Enthalpy is the total potential energy of a system. In many cases the only work “w” done on a system results in change in calorific value of the available mass. In other words: w ¼ DE þ 0 ¼ DE
(2.3)
Therefore, the heat absorbed (q) in a process, measured under conditions of constant volume, is equal to the internal energy increase. According to this equation, if no outside work is done, the energy absorbed by the system is equal to the potential internal increase. A biological example of this would be the reverse of the previous chemical reaction, or photosynthesis: 6 CO2 þ 6H2O þ 709 kcal / C6H12O6 þ 6 O2, {caloric difference ¼ change in vol. of CO2 versus O2} Enthalpy is defined as the heat content of a system. Bond energy is amount of energy required to break a chemical bond. The total bond energy is equivalent to the total potential energy of the system, a quantity known as enthalpy (H). The heat absorbed in a process at constant pressure is equal to the change in enthalpy, DH. Since a change in enthalpy can occur through both a change in pressure or volume, as well as internal energy, another term is introduced to describe the heat capacity of a substance: DH ¼ DE þ D (pv)
(2.4)
The specific heat of a substance is defined as the quantity of heat required to raise the temperature of 1 g of substance by 1 Celsius. This is an extremely important relationship for biological systems. It explains the importance of water as a “thermal buffer,” since when compared to other solvents water possesses relatively high heats of vaporization and freezing. The Second Law of Thermodynamics. We are all familiar with the fact that many energetic processes occur spontaneously. For example, water runs downhill; gases expand from regions of high pressure to regions of low pressure; chemical reactions proceed to equilibrium; and heat flows from warm bodies to cooler bodies. The Second Law of Thermodynamics states that processes involving transformations will not occur spontaneously, unless there is a degradation of energy from a nonrandom (ordered) form to a random
The physical and chemical bases of energy Chapter | 2
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(disordered) form. In natural systems, spontaneous energy transformations result in the degradation of the energy state of the system from a useful form to a dissipated and less usable form of heat. Obviously, as spontaneous processes occur in a system the system loses the ability to do work. Living (biological) systems have evolved to exploit these natural energy transformations and to utilize energy as it passes from ordered to random states.
All systems tend to approach states of equilibriumdin thermodynamic properties; this means complete randomness or energy degradation of the system. As a measure of the extent to which this equilibrium has been reached, another thermodynamic term, entropy is introduced. Entropy. Entropy is a measure of the disorder, or randomness, of a system. Organized, usable energy has low entropy, whereas disorganized entropy such as heat has high entropy. The more the molecules in a system are distributed in a disordered or random manner, the more probable is the arrangement and the greater is the entropy. The First Law of Thermodynamics recognizes the interconvertibility of all forms of energy, but it does not predict how complete the conversions will be. This applies to all energy conversions, except the transformation to heat, which is a property of molecules moving around at random. By contrast, all other forms of energy result from an ordered, nonrandom arrangement of the elementary particles of matter. Heat is the only from of energy due to disorder or random movement, and it is the most likely energy form to occur.
2.4 Gaia hypothesis The Gaia Hypothesis proposed by James Lovelock (1972) suggests that living organisms on the planet interact with their surrounding inorganic environment to form a synergetic and self-regulating system that created, and now maintains, the climate and biochemical conditions that make life on Earth possible. Gaia bases this postulate on the fact that the biosphere, and the evolution or organisms, affects the stability of global temperature, salinity of seawater, and other environmental variables. For instance, even though the luminosity of the sun, the Earth’s heat source, has increased about 30% since life began almost four billion years ago, the living system has reacted as a whole to maintain temperatures at a level suitable for life. Cloud formation over the open ocean is almost entirely a function of oceanic algae that emit sulfur molecules as waste metabolites which become condensation nuclei for rain. Clouds, in turn, help regulate surface temperatures. Lovelock compared the atmospheres of Mars and Earth, and noted that the Earth’s high levels of oxygen and nitrogen were abnormal and thermodynamically in disequilibrium. The 21% oxygen content of the atmosphere is an obvious consequence of living organisms, and the levels of other gases, NH3
12 The Global Carbon Cycle and Climate Change
and CH4, are higher than would be expected for an oxygen-rich atmosphere. Biological activity also explains why the atmosphere is not mainly CO2 and why the oceans are not more saline. Gaia postulates that conditions on Earth are so unusual that they could only result from the activity of the biosphere (Lovelock and Margulis, 1974).
2.5 Carbon and energy After the origin of the universe some 13e18 billion years ago with the Big Bang, a condensing sun began to collapse and increasing pressure allowed helium to “burn” to form carbon. 4
He þ 4He 4 8Be
(2.5)
Be þ 4He /
(2.6)
8
12
C
Hydrogen, H, and helium, He, were the original building blocks of the universe. Approximately 3.8 billion years ago when surface temperatures cooled to 100 C, water condensed out of the atmosphere to form the primitive oceans. Water vapor and carbon dioxidedboth degassing from the Earth’s crustdserved as an early greenhouse. This primordial atmosphere kept early Earth from freezing. Without the presence today of water vapor and CO2 in the atmosphere that creates a significant greenhouse effect, the Earth would be about 33 C cooler and covered by ice (Ramanathan, 1988). The forms of carbon. Carbon occurs in many different materials in many different forms. In bioenergetics, one is interested in the carbon content of organic molecules. Carbon content may be categorized as: Total Carbon (TC)dall the carbon in the sample, including both the inorganic and the organic carbon, Total Inorganic Carbon (TIC)d often referred to as inorganic carbon (IC), carbonate, bicarbonate, and dissolved carbon dioxide (CO2), Total Organic Carbon (TOC)dmaterial derived from decaying vegetation, bacterial growth, and metabolic activities of living organisms or chemicals, and Elemental Carbon (EC)dcharcoal, coal, and soot. Resistant to analytical digestion and extraction, EC can be a fraction of either TIC or TOC depending upon the analytical approach. Carbon mass may be calculated from the proportional composition of carbon in the substrate if its molecule composition is known, i.e., proportional mass of element ¼
mass of element in compound total mass of compound
(2.5)
Therefore, the proportion of carbon times the weight of the compound yields the mass of carbon present. Measurement of CO2, as well as O2 and CH4, gas
The physical and chemical bases of energy Chapter | 2
13
concentrations are typically made using gas chromatography and thermal conductivity (for example, Haskin, 2013; Emerson). Measures of carbon. In geochemistry, carbon is measured in units of mass, typically grams or kilograms (103 g). Molecular concentrations can be expressed in terms of mass: milligrams per liter (mg L1); or in terms of volume: microliters per liter (ul/L). These two expressions will be the same only if the density of particles is 1 g cm3. Atmospheric CO2 concentrations are given as ppm volume. Atmospheric concentrations of carbon in the form of carbon dioxide (CO2) are expressed as parts per million by volume (ppm volume) and can be compared relative to other molecular gases in the atmosphere. To convert from ppm by mass to ppm by volume, divide by the density of the molecules. Quantities may be expressed either as mass or as moles (mols). A mole is defined as the amount of a chemical substance that contains as many representative particles, i.e., atoms, as there are atoms in 12 g of 12C, which is Avogadro’s number ¼ 6.0221,409 1023. A mole of CO2 is one gram atomic weight, or the mass in grams of one mole of atoms; a mole of carbon dioxide is equivalent to [12 (C) þ 16 (O) x 2] ¼ 44g; while a mole of 12C is thus 12g C. Carbon chemistry. Organic molecules contain both carbon and hydrogen. Although many organic chemicals contain other elements, it is the carbonhydrogen bond that defines them as organic. Organic chemistry defines life. Just as there are millions of different types of living organisms on this planet, there are also millions of different kinds of organic molecules, each with different chemical and physical properties. Carbon (C) appears in the second row of the periodic table and has four bonding electrons in its valence shell. Similar to other nonmetals, carbon needs eight electrons to satisfy its valence shell. Carbon, therefore, forms four bonds with other atoms, each bond consisting of one of carbon’s electrons and one of the bonding atom’s electrons. Organic chemicals get their diversity from the many different ways carbon can bond to other atoms. The simplest organic chemicals, called hydrocarbons, contain only carbon and hydrogen atoms; the simplest hydrocarbon, methane (CH4), contains a single carbon atom bonded to four hydrogen atoms. To add to the complexity of organic chemistry, neighboring carbon atoms can form double and triple bonds in addition to single carbon-carbon bonds and join with other atoms, such as phosphorous and sulfur, forming complex chains and rings (Carpi, 2013; Hamilton, 2017). Stable chemical bonds release energy as they form, and bond formation thermodynamically happens spontaneously. However, formation reactions often require energy of activation to rearrange bonds and to get reactions over activation barriers, which are usually an exothermic breaking of bonds and the formation of new ones. Chemical bonds “contain” energy, but energy must be added to yield energy. In biochemical reactions, the energy for breaking bonds
14 The Global Carbon Cycle and Climate Change
comes from the formation of stronger bonds. In photosynthesis, energy from the sun breaks the CO2 and H2O bonds, and a fairly strong O2 bond is formed. Cellular respiration provides energy by forming the strong oxygen bonds in carbon dioxide and water, breaking the weaker bonds in carbohydrates and sugars. The greater the difference between the bond energies of the formed products (CO2 and H2O), and the reactants, the more energy becomes available. More energy is “available” when the weakest bonds are broken in favor of the stronger bonds being formed. For example, ATP, the coenzyme in cellular respiration, provides energy to chemical reactions in the metabolic process when it transfers phosphate moieties to more strongly bonded glucose or fructose phosphates (see Section 4.3 in Chapter 4).
2.6 Recommended reading Lovelock, J.E., Margulis, L., 1974. Atmospheric homeostasis by and for the biosphere: the Gaia hypothesis. Tellus 26 (1e2), 2e10. https://doi.org/10.3402/tellusa.v26i1-2.9731. Morowitz, H.J., 1970. Entrophy for Biologists. Academic Press, 195 pp. https://books.google.com/ books?hl=en&lr=&id=JSrLBAAAQBAJ&oi=fnd&pg=PP1&ots=QaAM4BSjNz&sig= jFAYu5hagX6HTST-xSCq1094Mbk#v=onepage&q&f=false. Wikipedia, 2019. Laws of Thermodynamics. https://en.wikipedia.org/w/index.php?title¼Laws_ of_thermodynamics&oldid¼888134088/.
The physical and chemical bases of energy 2.1 Energy, work, and power Calories and Joules 2.2 The different forms of energy Chemical energy Radiant energy Heat energy 2.3 The Laws of Thermodynamics The First Law of Thermodynamics Work Enthalpy The Second Law of Thermodynamics Entropy 2.4 Gaia hypothesis 2.5 Carbon and energy The forms of carbon Measures of carbon Carbon chemistry 2.6 Recommended reading
There is no better way to begin the study of ecological energetics than by starting with an understanding of the pertinent definitions and terminology of physics and physical chemistry. Learn this terminology early, become comfortable with the units of measure, know the basic concepts, and bioenergetics will come a lot easier.
2.1 Energy, work, and power Energy is the capacity to do work. The unit of measure for energy is the erg, which is the work performed when a force of one dyne acts through a distance of one centimeter. The unit of force, the dyne, yields to a mass of one gram the acceleration of one centimeter per second (cm s1). Since an erg of energy is such a small quantity, a larger unit, the joule, which is equal to 107 ergs, becomes a more convenient unit of measure. A unit of heat used frequently in The Global Carbon Cycle and Climate Change. https://doi.org/10.1016/B978-0-12-820244-9.00002-0 Copyright © 2020 Elsevier Inc. All rights reserved.
5
6 The Global Carbon Cycle and Climate Change
physical chemistry is the calorie (¼ 4.184 J). The calorie is the heat energy required to raise the temperature of one gram of water from 14.5 oC to 15.5oC. The calorie is defined as being equal to 4.1840 absolute joules. The calorie is a relatively small unit of measure, and for most chemical and biological calculations the kilocalorie (103 calories) is used. The kilocalorie (kcal) is the unit which is typically used in discussing dietary intake and is often written as Calorie. A Calorie equals 103 calories, or a kcal. Calories and Joules. A calorie is the energy needed to raise the temperature of 1 g of water through 1 C (also expressed as 4.1868 J, the unit of energy in the International System of Units). A joule is the energy expended when 1 kg is moved 1 m by a force of 1 Newton (N). Use of joules is now recommended by international convention and is the preferred standard unit to measure heat (FAO, 2003). Nutritionists and food scientists concerned with large amounts of energy generally use kiloJoules (kJ ¼ 103 J) or megaJoules (MJ ¼ 106 J). For many decades, food energy has been expressed in calories, and studies in the field of ecological energetics have traditionally used calories as the measure for energy. In order to retain consistency with research reported in the scientific literature, values used for energy in this book are in calories. The conversion factors for joules and calories are: 1 cal ¼ 4.184 J and 1 J ¼ 0.239 cal.
2.2 The different forms of energy Energy can exist in various forms, but those of greatest importance to living organisms are mechanical, chemical, radiant, and heat energy (Table 2.1). Mechanical energy has two forms: kinetic and potential. Kinetic energy, or free energy, can be described as the “useful energy” which a body possesses by
TABLE 2.1 Units of measure for energy in its various forms and transformations. Energy
Intensity
Mechanical (ergs)
Force (dynes)
Kinetic (ergs) Potential (ergs) Chemical (calories)
Capacity Change in distance (cm) -1
1
Velocity (cm s )
/2 Mass (g)
-2
Height x accereraltion (cm s ) 1
Heat of combustion (cal g ) 2
Mass (g) Mass (g)
Radiant (calories)
Radiation flux (cal cm )
Surface area (cm2)
Heat (calories)
Difference in temperature (o C)
Heat capacity (cal per oC)
Vol. expansion (ergs)
Pressure (dynes cm-2)
Change in volume (cm3)
Electrical (joules)
Difference in potential (volts)
Surface (ergs)
1
Surface tension (ergs cm )
Coulombs (amps x sec.) Change in area (cm2)
The physical and chemical bases of energy Chapter | 2
7
value of its motion, and is measured by the amount of work which is done in bringing that body to rest. Examples would be a moving ball or the Brownian movement of molecules. Potential energy is stored energy, which is only potentially useful until its conversion into the kinetic or free energy where it becomes available to accomplish work. Energy may be stored in a system by virtue of position, as for example, a stone above the Earth’s surface, a steel spring under compression, or by virtue of chemical properties due to the arrangement of atoms and electrons within a molecule. Conversion of energy from the potential form to the kinetic form involves movement, i.e., motion. Chemical energy. All organisms must work to live, and they require a source of potential energy which can be utilized in order to perform the life processes. This energy can be found in the form of the chemical energy of biomass used as food. Energy can also be in the form of the chemical energy of inorganic molecules utilized as an alternative energy source to radiant energy by chemotrophs. Assemblies of atoms in matter can be rearranged into different groups; thus, by the movement of atoms and the creation of different atomic bonds, chemical energy is liberated. The combustion (oxidation) of coal in a furnace or food by the respiratory processes in a cell releases energy which can be used to accomplish work. Both of these processes illustrate the conversion of chemical to mechanical energy. Life processes on this Earth have evolved around carbon chemistry, and most chemical energy sources are derived from organic compounds. However, as we shall see shortly, there are some notable exceptions. Radiant energy. The sun, a vast incandescent sphere of gas, releases energy by the nuclear transmutation of hydrogen to helium, and it is upon this energy source that life on Earth depends. Radiant energy is the energy of electromagnetic radiation. Because electromagnetic radiation can be conceptualized as a stream of photons, radiant energy can be viewed as photon energy. Alternatively, EM radiation can be viewed as an electromagnetic wave, which carries energy in its oscillating electric and magnetic fields. These two views are completely equivalent and are reconciled to one another in quantum field theory. Solar radiation is energy in the form of electromagnetic waves involving a rhythmic exchange between potential and kinetic energy. Electromagnetic radiation can have frequencies, or wave lengths, of different energy content and interactions (e.g., absorptivities) with matter. Heat energy. This is a very special form of energy resulting from the random movements of molecules, which by virtue of their motion, possess kinetic energy. Heat is evolved when all other forms of energy are transformed and work is performed. All work, including the growth and reproduction of living organisms, represent the transformation of energy and ultimately results in the production of heat. For example, when an animal during respiration releases the potential energy of glucose, approximately two-thirds of it is converted
8 The Global Carbon Cycle and Climate Change
into mechanical energy to be used for work (activity and growth) and onethirds is given off as heat. There are instances of work where heat is absorbed (endodermic processes): the cooling unit of a refrigerator or the fixation of atmospheric nitrogen by certain bacteria are examples; but, these processes are not self-supporting energetically. Nitrogen fixation is always accompanied by the exothermic breakdown of organic substrates. Heat energy released by an exothermic process is never used with complete efficiency by the endergonic process, and so whenever work is done the trend is always toward heat production. In natural processes, changes from one form of energy to another (except to heat) are normally incomplete, because the movement, already shown to be necessary for energy conversion, involves either friction or heat production. Temperature is the relative measure which is used to characterize the amount of heat present in a system. Several common systems are used, only two of which concern us: the Celsius and the Kelvin scales. The Celsius (centigrade) scale (oC) establishes zero (0 ) as the freezing point of water and 100 degrees as the boiling point of water. The Kelvin scale is an absolute measure which establishes zero at the temperature (273 C) at which all molecular motion ceases. Because freezing and boiling points limit life processes, we will conventionally utilize the Celsius scale. Heat may be characterized by the properties of two phases: sensible heat and latent heat. The significance of each of these phases will shortly become apparent. Suffice at present to distinguish between the two phases as follows: sensible heat is that which can be measured by an increase in temperature of a body, for example, the warming action of sunlight irradiating a metal plate. Latent heat is the heat absorption by a body without an equivalent increase in temperature, such as the heats of freezing or vaporization of water (heat of vaporization of water ¼ 539 cal at 100 C). Energy flow is expressed as the product of two factors: (1) an intensity factor (or gradient), and (2) a capacity factor (amount). Energy, work, and heat are all expressed in the same units: calories, joules, or ergs. It should be evident that the different energies may be compared, but that no relationship exists between the capacity factors alone. For example, electric energy may be converted into heat energy, but the rise in temperature cannot be calculated from the voltage, unless the number of coulombs and the heat capacity of the system are known. It is also clear that the same quantity of work can be accomplished by a small quantity of water passing through a turbine from a great height as by a large quantity of water passing through from a short distance. Thus, arises one of the fundamental principles of thermodynamicsdthe interconvertibility of energy forms as well as the “trade-offs” between the intensity and capacity factors. All forms of energy are interconvertible; when conversions do occur, they do so according to rigorous laws of exchange. These are the Laws of Thermodynamics.
The physical and chemical bases of energy Chapter | 2
9
2.3 The Laws of Thermodynamics The First Law of Thermodynamics. The First Law of Thermodynamics is also known as the Law of Conservation of Energy, since it defines that the sum of all energies in an isolated system is constant. In other words, energy may be transformed from one state to another, but it can neither be created nor destroyed. The total energy in universe remains constant, but it is continuously becoming more diffuse throughout the universe. Remember that the capacity factor for mechanical and chemical energy is mass (m). Thus, Einstein showed that if there is a change in mass, Dm, Energy ¼ Dmc2,
(2.1)
where: c is the velocity of light (3 1010 cm s1). Therefore, 1 g of water is equivalent to 9 1020 ergs of energy. It should be evident that the Law of Conservation of Energy and the Law of Conservation of Mass are essentially the same, and that no violation of thermodynamics occurs when energy is converted into mass or mass is converted into energy. However, a change in energy of a system will be brought about if the system does work, or if it absorbs or evolves heat. Thus, when a change of any kind occurs in a closed system (where the amount of matter is fixed but energy is able to enter or leave) an increase or decrease occurs in the internal energy (E) of the system itself; heat (q) is evolved or absorbed, and work (w) is done: DE
the decrease in internal energy ðcalÞ
¼
q the heat ðcalÞ given off by the system
w
the work ð-calÞ done by the system
(2.2)
Work. The work performed by a system is the energy transferred by the system to its surroundings. The negative value of work indicates that a positive amount of work done by the system has led to energy being lost from the system. The First Law of Thermodynamics also encompasses the more specific relationship of constant heat sums, which is of considerable importance to biologists interested in energy transformations. It states that the total amount of heat produced, or absorbed, from a chemical reaction which takes place in stages is equal to the total amount of heat evolved, or consumed, when the reaction occurs directly in one step. The evolution of living systems has utilized biochemical mechanisms by which chemical compounds can be reduced in steps, thus enabling more efficient energy capture and utilization to occur. A good biological example is the metabolic oxidation of glucose to carbon dioxide and water: Direct reaction (combustion) C6H12O6 þ 6O2 /6H2O þ 6CO2 þ 673 kcal of energy
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Two-stage reaction (fermentation) (a) C6H12O6 /2C2H5OH þ 2CO2 þ 18 kcal of energy (b) 2C2H5OH þ 6O2 /6H2O þ 4CO2 þ 655 kcal of energy (a) þ (b) C6H12O6 þ 6O2 /6H2O þ 6CO2 þ 673 kcal of energy Thus, no matter which pathway a particular reaction follows, the total amount of heat evolved, or absorbed, is always the same. There will be more discussion about the biological significance of this phenomenon later. Several other energy relationships are also pertinent to bioenergetics. Enthalpy. Enthalpy is the total potential energy of a system. In many cases the only work “w” done on a system results in change in calorific value of the available mass. In other words: w ¼ DE þ 0 ¼ DE
(2.3)
Therefore, the heat absorbed (q) in a process, measured under conditions of constant volume, is equal to the internal energy increase. According to this equation, if no outside work is done, the energy absorbed by the system is equal to the potential internal increase. A biological example of this would be the reverse of the previous chemical reaction, or photosynthesis: 6 CO2 þ 6H2O þ 709 kcal / C6H12O6 þ 6 O2, {caloric difference ¼ change in vol. of CO2 versus O2} Enthalpy is defined as the heat content of a system. Bond energy is amount of energy required to break a chemical bond. The total bond energy is equivalent to the total potential energy of the system, a quantity known as enthalpy (H). The heat absorbed in a process at constant pressure is equal to the change in enthalpy, DH. Since a change in enthalpy can occur through both a change in pressure or volume, as well as internal energy, another term is introduced to describe the heat capacity of a substance: DH ¼ DE þ D (pv)
(2.4)
The specific heat of a substance is defined as the quantity of heat required to raise the temperature of 1 g of substance by 1 Celsius. This is an extremely important relationship for biological systems. It explains the importance of water as a “thermal buffer,” since when compared to other solvents water possesses relatively high heats of vaporization and freezing. The Second Law of Thermodynamics. We are all familiar with the fact that many energetic processes occur spontaneously. For example, water runs downhill; gases expand from regions of high pressure to regions of low pressure; chemical reactions proceed to equilibrium; and heat flows from warm bodies to cooler bodies. The Second Law of Thermodynamics states that processes involving transformations will not occur spontaneously, unless there is a degradation of energy from a nonrandom (ordered) form to a random
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(disordered) form. In natural systems, spontaneous energy transformations result in the degradation of the energy state of the system from a useful form to a dissipated and less usable form of heat. Obviously, as spontaneous processes occur in a system the system loses the ability to do work. Living (biological) systems have evolved to exploit these natural energy transformations and to utilize energy as it passes from ordered to random states.
All systems tend to approach states of equilibriumdin thermodynamic properties; this means complete randomness or energy degradation of the system. As a measure of the extent to which this equilibrium has been reached, another thermodynamic term, entropy is introduced. Entropy. Entropy is a measure of the disorder, or randomness, of a system. Organized, usable energy has low entropy, whereas disorganized entropy such as heat has high entropy. The more the molecules in a system are distributed in a disordered or random manner, the more probable is the arrangement and the greater is the entropy. The First Law of Thermodynamics recognizes the interconvertibility of all forms of energy, but it does not predict how complete the conversions will be. This applies to all energy conversions, except the transformation to heat, which is a property of molecules moving around at random. By contrast, all other forms of energy result from an ordered, nonrandom arrangement of the elementary particles of matter. Heat is the only from of energy due to disorder or random movement, and it is the most likely energy form to occur.
2.4 Gaia hypothesis The Gaia Hypothesis proposed by James Lovelock (1972) suggests that living organisms on the planet interact with their surrounding inorganic environment to form a synergetic and self-regulating system that created, and now maintains, the climate and biochemical conditions that make life on Earth possible. Gaia bases this postulate on the fact that the biosphere, and the evolution or organisms, affects the stability of global temperature, salinity of seawater, and other environmental variables. For instance, even though the luminosity of the sun, the Earth’s heat source, has increased about 30% since life began almost four billion years ago, the living system has reacted as a whole to maintain temperatures at a level suitable for life. Cloud formation over the open ocean is almost entirely a function of oceanic algae that emit sulfur molecules as waste metabolites which become condensation nuclei for rain. Clouds, in turn, help regulate surface temperatures. Lovelock compared the atmospheres of Mars and Earth, and noted that the Earth’s high levels of oxygen and nitrogen were abnormal and thermodynamically in disequilibrium. The 21% oxygen content of the atmosphere is an obvious consequence of living organisms, and the levels of other gases, NH3
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and CH4, are higher than would be expected for an oxygen-rich atmosphere. Biological activity also explains why the atmosphere is not mainly CO2 and why the oceans are not more saline. Gaia postulates that conditions on Earth are so unusual that they could only result from the activity of the biosphere (Lovelock and Margulis, 1974).
2.5 Carbon and energy After the origin of the universe some 13e18 billion years ago with the Big Bang, a condensing sun began to collapse and increasing pressure allowed helium to “burn” to form carbon. 4
He þ 4He 4 8Be
(2.5)
Be þ 4He /
(2.6)
8
12
C
Hydrogen, H, and helium, He, were the original building blocks of the universe. Approximately 3.8 billion years ago when surface temperatures cooled to 100 C, water condensed out of the atmosphere to form the primitive oceans. Water vapor and carbon dioxidedboth degassing from the Earth’s crustdserved as an early greenhouse. This primordial atmosphere kept early Earth from freezing. Without the presence today of water vapor and CO2 in the atmosphere that creates a significant greenhouse effect, the Earth would be about 33 C cooler and covered by ice (Ramanathan, 1988). The forms of carbon. Carbon occurs in many different materials in many different forms. In bioenergetics, one is interested in the carbon content of organic molecules. Carbon content may be categorized as: Total Carbon (TC)dall the carbon in the sample, including both the inorganic and the organic carbon, Total Inorganic Carbon (TIC)d often referred to as inorganic carbon (IC), carbonate, bicarbonate, and dissolved carbon dioxide (CO2), Total Organic Carbon (TOC)dmaterial derived from decaying vegetation, bacterial growth, and metabolic activities of living organisms or chemicals, and Elemental Carbon (EC)dcharcoal, coal, and soot. Resistant to analytical digestion and extraction, EC can be a fraction of either TIC or TOC depending upon the analytical approach. Carbon mass may be calculated from the proportional composition of carbon in the substrate if its molecule composition is known, i.e., proportional mass of element ¼
mass of element in compound total mass of compound
(2.5)
Therefore, the proportion of carbon times the weight of the compound yields the mass of carbon present. Measurement of CO2, as well as O2 and CH4, gas
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concentrations are typically made using gas chromatography and thermal conductivity (for example, Haskin, 2013; Emerson). Measures of carbon. In geochemistry, carbon is measured in units of mass, typically grams or kilograms (103 g). Molecular concentrations can be expressed in terms of mass: milligrams per liter (mg L1); or in terms of volume: microliters per liter (ul/L). These two expressions will be the same only if the density of particles is 1 g cm3. Atmospheric CO2 concentrations are given as ppm volume. Atmospheric concentrations of carbon in the form of carbon dioxide (CO2) are expressed as parts per million by volume (ppm volume) and can be compared relative to other molecular gases in the atmosphere. To convert from ppm by mass to ppm by volume, divide by the density of the molecules. Quantities may be expressed either as mass or as moles (mols). A mole is defined as the amount of a chemical substance that contains as many representative particles, i.e., atoms, as there are atoms in 12 g of 12C, which is Avogadro’s number ¼ 6.0221,409 1023. A mole of CO2 is one gram atomic weight, or the mass in grams of one mole of atoms; a mole of carbon dioxide is equivalent to [12 (C) þ 16 (O) x 2] ¼ 44g; while a mole of 12C is thus 12g C. Carbon chemistry. Organic molecules contain both carbon and hydrogen. Although many organic chemicals contain other elements, it is the carbonhydrogen bond that defines them as organic. Organic chemistry defines life. Just as there are millions of different types of living organisms on this planet, there are also millions of different kinds of organic molecules, each with different chemical and physical properties. Carbon (C) appears in the second row of the periodic table and has four bonding electrons in its valence shell. Similar to other nonmetals, carbon needs eight electrons to satisfy its valence shell. Carbon, therefore, forms four bonds with other atoms, each bond consisting of one of carbon’s electrons and one of the bonding atom’s electrons. Organic chemicals get their diversity from the many different ways carbon can bond to other atoms. The simplest organic chemicals, called hydrocarbons, contain only carbon and hydrogen atoms; the simplest hydrocarbon, methane (CH4), contains a single carbon atom bonded to four hydrogen atoms. To add to the complexity of organic chemistry, neighboring carbon atoms can form double and triple bonds in addition to single carbon-carbon bonds and join with other atoms, such as phosphorous and sulfur, forming complex chains and rings (Carpi, 2013; Hamilton, 2017). Stable chemical bonds release energy as they form, and bond formation thermodynamically happens spontaneously. However, formation reactions often require energy of activation to rearrange bonds and to get reactions over activation barriers, which are usually an exothermic breaking of bonds and the formation of new ones. Chemical bonds “contain” energy, but energy must be added to yield energy. In biochemical reactions, the energy for breaking bonds
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comes from the formation of stronger bonds. In photosynthesis, energy from the sun breaks the CO2 and H2O bonds, and a fairly strong O2 bond is formed. Cellular respiration provides energy by forming the strong oxygen bonds in carbon dioxide and water, breaking the weaker bonds in carbohydrates and sugars. The greater the difference between the bond energies of the formed products (CO2 and H2O), and the reactants, the more energy becomes available. More energy is “available” when the weakest bonds are broken in favor of the stronger bonds being formed. For example, ATP, the coenzyme in cellular respiration, provides energy to chemical reactions in the metabolic process when it transfers phosphate moieties to more strongly bonded glucose or fructose phosphates (see Section 4.3 in Chapter 4).
2.6 Recommended reading Lovelock, J.E., Margulis, L., 1974. Atmospheric homeostasis by and for the biosphere: the Gaia hypothesis. Tellus 26 (1e2), 2e10. https://doi.org/10.3402/tellusa.v26i1-2.9731. Morowitz, H.J., 1970. Entrophy for Biologists. Academic Press, 195 pp. https://books.google.com/ books?hl=en&lr=&id=JSrLBAAAQBAJ&oi=fnd&pg=PP1&ots=QaAM4BSjNz&sig= jFAYu5hagX6HTST-xSCq1094Mbk#v=onepage&q&f=false. Wikipedia, 2019. Laws of Thermodynamics. https://en.wikipedia.org/w/index.php?title¼Laws_ of_thermodynamics&oldid¼888134088/.