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Chapter 6: The basics of chemistry and interaction of electromagnetic radiation (light and heat) with matter. 6.1 Basic chemical concepts Protons, neutrons, electrons, and electrostatic forces. Atoms are the fundamental chemical building blocks of matter, the smallest unit that retains chemical identity. An atom is made up of protons (positive charge), neutrons (zero charge), and electrons (negative charge). The protons and neutrons are packed together in the nucleus and the electrons forming a cloud of negative charge around the nucleus. The extensive distribution of electrons around the nucleus is a consequence of the fact that the mass of an electron is only 1/2000th that of a proton or neutron. The typical size of an atom (diameter of the electron cloud) is 10-10 m, whereas the nucleus is smaller by a factor of 100. The atomic unit of length is the Ångström, 1 Å ≡ 10-10 m, named to honor Anders Ångström (1814-1874), Swedish physicist who improved the precision in measuring the wavelength of spectral lines. Electrostatic forces are responsible for holding atoms together or forcing them apart. When electric charges, q1 and q2 are distance r apart, the electrostatic force between them is F = q1q2/r2, and the energy of interaction (the energy it takes to bring them to distance r from infinity) is E = q1q2/r . If the charges are of like sign (both + or both -), the force is repulsive, the energy is positive, and the charges will tend to fly apart. Charges with opposite signs are attracted and the electrostatic force pulls the charges together. When atoms approach each other, the clouds of electrons begin to interact when the electrons start to overlap. You might expect that the like charges on the electrons would repel all approaching atoms. This is often, but not always, true. Under some circumstances, the close approach of two atoms causes the outermost electrons on both to rearrange themselves, following complex rules described by the laws of quantum physics. Figure 6.1 shows an example of the distribution of charge in a molecule, Si3. Figure 6.1 Illustration showing calculated density of electron charge (net negative charge, shown in red and green) relative to the positions of the nuclei and inner-shell electrons (net positive charge, dark blue) in the molecule Si3. The maxima of electron density between the nuclei provide clouds of negative charge that attract the positively-charge nuclei and hold the molecule together. (Figure by Dr. Masao Arai, National Institute for Research in Inorganic Materials, Japan; reproduced by permission..) 1 Note how some electrons have bunched together between the nuclei. The nuclei (positive charges) are attracted to the electrons bunched between them, providing an attractive force. Figure 6.1 illustrates a chemical bond between the atoms, holding them together in a molecule. Elements, isotopes, shells of electrons and the Periodic Table The mass of an atom is determined essentially by the mass number, the number of protons + the number of neutrons, since electrons have very small mass. However, the number of electrons in an atom determine the rules for interaction with the electrons of another atom, and hence its chemical properties. Thus the chemical properties of a neutral atom are determined by its atomic number, the number of protons in the nucleus, and an element is defined as atom with given atomic number. A mole is N0 = 6.023 1023 atoms (Avogadro's number). The mass (in grams) of one mole equals the mass number of an element. For example, the mass of one mole of 12C is 12 grams. Similarly, a mole of a chemical compound (molecule with more than one atom) represents 6.023 1023 molecules and has a mass in grams equal to the sum of mass numbers of its atoms. The number of neutrons in an atom of a given element may vary. The sum of protons and neutrons is the mass number of the atom. For example, the atom 12C has 6 neutrons and 6 protons so its mass number is 12. Atoms of the same element but with different numbers of neutrons are called isotopes. Isotopes of carbon are 12C, 13C, 14 C. These atoms have 6, 7, and 8 neutrons respectively, but all have 6 protons and 6 electrons. Since isotopes of an element all have the same number of electrons, they have essentially the same behavior when forming molecules. Slight differences in the chemical behavior of isotopes can provide important information about molecules found in nature. Elements form chemical compounds in very systematic ways, following patterns recognized long before the true nature of atoms was understood. The 19th century Russian chemist Dmitri Mendeleev codified these patterns in the periodic table, shown below. Elements with similar chemical behavior are arranged in columns in order of increasing mass for the atom (atomic weights were already known for many elements in Mendeleev's time). The sequential filling up of each row of the table follows a progression with increasing atomic mass, starting with metallic elements, progressing to non-metals that can attack metals to make salts, and finally "noble gases" that will not bond with other atoms at all. Then the pattern starts over with the next heavier element. The orderly pattern of chemical bonding reflects the fact that electrons are arranged in groupings called shells, with all electrons in pairs. The most stable configurations for electron shells correspond to special numbers: 2 (1 pair), 8 (4 pairs), and 8 (4 pairs), for rows 1, 2 , 3 respectively. Each noble gas has exactly such a number of electrons in its outermost shell, called a closed shell, corresponding to atomic number 2, 10 (2+8), and 18 (2+8+8). The progression along each row of the periodic table represents the filling up of the outermost shell until the stable configuration is attained. 2 The metals, on the left side of each row, have one or two electrons in the outer shell and most readily attain a closed shell by giving up these outermost electrons. For example, the sodium atom has just one electron in its outer shell, and the sodium ion Na+ has a closed shell. Electrons move easily from one atom to another, allowing metals to conduct electric currents. The non-metal chorine, on the right side of the row, needs to acquire a single electron, becoming Cl–, to obtain a closed shell configuration. Thus sodium metal and chlorine gas are both quite unstable and reactive, but table salt, NaCl, has the form Na+ Cl– and is very stable. If the numbers of electrons and protons are unequal in an atom or molecule, it is called an ion. Cations are ions with net positive charge (fewer electrons than protons), anions have net negative charge (more electrons than protons). Molecules dissolved in water often dissociate to form ions, because water molecules attach to these ions and make them more stable than non-ionized species. For example, hydrochloric acid (HCl) dissociates to H+ and Cl– ions, each surrounded by water molecules (Fig. 6.2). The acidity of a water solution is defined by the concentration of H+ ions, measured in units of moles per liter of solution. A common unit of acidity is pH ≡ -log10[H+], where [X] denotes the number of moles of X per liter of solution. Figure 6.2. Water molecules forming solvent shells around aqueous ions. The O atom in water partially pulls the electrons away from the H atoms, giving its side of the molecule a small negative charge (-2δ) and the H side a small positive charge (+δ on each H-atom). Molecules of water (the solvent) cluster around ions in solution, with the opposite charge pointing towards the ion, to form a shell. The surrounding shell of water molecules stabilizes the ions in solution, making the charges on other ions "invisible" and keeping them from rejoining each other. Most stable molecules are formed when each of the atoms is able to attain a closed shell by sharing electrons with adjacent atoms to make a covalent bond , and in a few cases the electrons are completely exchanged, as in the ionic bond of the molecule NaCl. Examples include H2O (water), CO2 (carbon dioxide), CH4 (methane) (see Fig. 6.3). Only the outermost electrons are involved in the chemical bond. There are inner shells of electrons tightly bound around the nuclei. For example, in Figure 6.1 these electrons are shown near the nuclei for the molecule Si3. Since the number of electrons over all equals the charge on the nuclei, and the charge of (nucleus + inner electrons) is positive. These complexes of (nucleus + inner electrons) are held together by electrostatic forces to make a molecule. 3 A useful way to visualize the sharing of electrons to create covalent bonds is illustrated in Figure 6.3. The electrons in the outer (valence) shell are depicted as dots, and inner-shell electrons are ignored. For the 2nd-row elements O and C, eight electrons make up the outer shell, and for the H atoms, two are needed. To make the water molecule, the Oatom, with six electrons, shares one pair of electrons with each H atom, and all the atoms attain closed shells. A more complex sharing arrangement is needed for the CO2 molecule. Each O atom is looking to share two pairs of electrons in order to pick up the two that it is missing, and the C atom, initially with four valence electrons, needs to share four pairs to pick up four more. This is accomplished as depicted in Figure 6.3. Each bond between C and O involves two pairs of shared electrons, and is therefore double bond. As one might expect, double bonds are stronger than single bonds. Figure 6.3 Valence electrons and chemical bonds. These diagrams illustrate the number of electrons in the outermost shell, called the valence shell, for the common atoms hydrogen, carbon, and oxygen. The electrons are shared when the atoms form stable compounds, so that every atom in the molecule has the magic number of electrons in its outer shell, giving it a closed shell (2 for H, 8 for C and O). Example: Find the valence bond structure for the following atmospheric gases: (a) molecular nitrogen (N2), (b) methane (CH4), and (c) nitric oxide (NO). (a) The N atom has five electrons in its outer shell. Each N atom needs to acquire three electrons by sharing. Each shared pair contributes one, and the nitrogen molecule has a bond in which three pairs of electrons are shared. This very strong triple bond accounts for the stability of the N2 molecule. (b) The carbon atom has four electrons in its outer shell, and each hydrogen atom has one. The carbon atom achieves a closed shell by acquiring four electrons, sharing one pair with each of the the four H atoms. (c) The N atom has five electrons in its outer shell and the O atom has six. Nitric oxide has an odd number of electrons, and there is no way have all the electorns paired in order to obtain a closed shell configuration. One remains unpaired, and therefore nitric oxide is a free radical that reacts rapidly with a wide variety of other gases. 4 WebElements: the periodic table on the world-wide web http://www.shef.ac.uk/chemistry/web-elements/ 1 hydrogen 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 1 2 H 1.00794(7) lithium 3 Li 6.941(2) sodium 11 He Key: beryllium 4 element name boron atomic number Be 5 B element symbol 9.012182(3) magnesium 1995 atomic weight (mean relative mass) 10.811(7) aluminium 12 13 Na Mg 22.989770(2) potassium 19 K 39.0983(1) rubidium 37 Al 24.3050(6) calcium scandium 20 21 Ca Sc 40.078(4) strontium 44.955910(8) yttrium 38 39 Rb Sr 85.4678(3) caesium 55 87.62(1) barium 56 Cs Ba 132.90545(2) francium 87 137.327(7) radium 88 Fr Ra [223.0197] 18 helium [226.0254] Y 57-70 * 89-102 ** lanthanum 57 *lanthanides 71 Lu 174.967(1) lawrencium 103 Lr [262.110] cerium 58 22 Ti 47.867(1) zirconium 40 vanadium 23 V 50.9415(1) niobium 41 chromium 24 manganese 25 iron 26 cobalt 27 nickel 28 copper 29 zinc 30 72 92.90638(2) tantalum 73 Hf Ta 178.49(2) rutherfordium 104 180.9479(1) dubnium 105 C 12.0107(8) silicon 14 Si 28.0855(3) germanium 32 nitrogen 7 N 14.00674(7) phosphorus 15 P 30.973761(2) arsenic 33 oxygen 8 O 15.9994(3) sulfur 16 S 32.066(6) selenium 34 fluorine 9 F 18.9984032(5) chlorine 17 Cl 35.4527(9) bromine 35 Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br 51.9961(6) 54.938049(9) molybdenum technetium 42 43 95.94(1) tungsten [98.9063] rhenium 55.845(2) ruthenium 44 58.933200(9) rhodium 45 58.6934(2) palladium 46 63.546(3) silver 47 65.39(2) cadmium 48 Zr Nb Mo Tc Ru Rh Pd Ag Cd 91.224(2) hafnium 31 6 74 75 101.07(2) osmium 76 W Re Os 183.84(1) seaborgium 106 186.207(1) bohrium 107 190.23(3) hassium 108 102.90550(2) iridium 77 Ir 192.217(3) meitnerium 109 106.42(1) platinum 78 107.8682(2) gold 79 112.411(8) mercury 80 Pt Au Hg 195.078(2) ununnilium 110 196.96655(2) unununium 111 200.59(2) ununbium 69.723(1) indium 49 In 72.61(2) tin 50 74.92160(2) antimony 51 118.710(7) lead Tl Pb 207.2(1) 208.98038(2) erbium thulium ytterbium 204.3833(2) 52 Sn Sb Te 114.818(3) thallium 81 78.96(3) tellurium 82 121.760(1) bismuth 83 Bi 127.60(3) polonium 84 79.904(1) iodine 53 I 126.90447(3) astatine 85 10 Ne 20.1797(6) argon 18 Ar 39.948(1) krypton 36 Kr 83.80(1) xenon 54 Xe 131.29(2) radon 86 Po At Rn [208.9824] [209.9871] [222.0176] 112 Rf Db Sg Bh Hs Mt Uun Uuu Uub [261.1089] [262.1144] [263.1186] [264.12] [265.1306] [268] [269] [272] [277] praseodymium neodymium promethium samarium europium gadolinium terbium dysprosium holmium 59 60 140.90765(2) protactinium 144.24(3) uranium 61 62 63 64 65 66 67 68 69 70 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb 138.9055(2) actinium 89 **actinides 88.90585(2) lutetium titanium 26.981538(2) gallium carbon 4.002602(2) neon 140.116(1) thorium 90 91 Ac Th Pa [227.0277] 232.0381(1) 231.03588(2) 92 U 238.0289(1) [144.9127] neptunium 93 150.36(3) plutonium 94 151.964(1) americium 95 157.25(3) curium 96 158.92534(2) berkelium 97 162.50(3) californium 98 164.93032(2) einsteinium 99 167.26(3) fermium 100 168.93421(2) mendelevium 101 173.04(3) nobelium 102 Np Pu Am Cm Bk Cf Es Fm Md No [237.0482] [244.0642] [243.0614] [247.0703] [247.0703] [251.0796] [252.0830] [257.0951] [258.0984] [259.1011] 5 Symbols and names: the symbols of the elements, their names, and their spellings are those recommended by IUPAC. After some controversy, the names of elements 101-109 are now confirmed: see Pure & Appl. Chem., 1997, 69, 2471Ð2473. Names have not been proposed as yet for the most recently discovered elements 110Ð112 so those used here are IUPACÕs temporary systematic names: see Pure & Appl. Chem., 1979, 51, 381Ð384. In the USA and some other countries, the spellings aluminum and cesium are normal while in the UK and elsewhere the usual spelling is sulphur. Periodic table organisation: for a justification of the positions of the elements La, Ac, Lu, and Lr in the WebElements periodic table see W.B. Jensen, ÒThe positions of lanthanum (actinium) and lutetium (lawrencium) in the periodic tableÓ, J. Chem. Ed., 1982, 59, 634Ð636. Group labels: the numeric system (1Ð18) used here is the current IUPAC convention. For a discussion of this and other common systems see: W.C. Fernelius and W.H. Powell, ÒConfusion in the periodic table of the elementsÓ, J. Chem. Ed., 1982, 59, 504Ð508. Atomic weights (mean relative masses): see Pure & Appl. Chem., 1996, 68, 2339Ð2359. These are the IUPAC 1995 values. Elements for which the atomic weight is contained within square brackets have no stable nuclides and are represented by one of the elementÕs more important isotopes. However, the three elements thorium, protactinium, and uranium do have characteristic terrestrial abundances and these are the values quoted. The last significant figure of each value is considered reliable to ±1 except where a larger uncertainty is given in parentheses. ©1998 Dr Mark J Winter [University of Sheffield, [email protected]]. For updates to this table see http://www.shef.ac.uk/chemistry/web-elements/pdf/periodic-table.html. Version date: 1 March 1998. 6.2 Chemical reactions and equilibrium Chemical reactions can occur by rearranging the bonds between atoms. If the rearrangement lowers the energy of the molecules, then energy is released to the environment, either as heat or as some other form of chemical energy, and the reaction is said to be exothermic. Reactions that require energy from the environment in order to proceed are called endothermic. For example, the reaction 2 H2 + O2 → 2 H2O is exothermic, with XXX Joules released for each mole of H2O produced; in fact an explosion or fire results when hydrogen and oxygen are mixed together and ignited. The reverse reaction, 2 H2O → 2 H2 + O2, occurs only at high temperatures, when there is abundant heat from the environment to provide the required energy. It is a curious fact that reactions that can proceed often do not. The rearrangements of electrons required to break existing bonds and make new ones are too improbable to actually occur. For example, when hydrogen and oxygen are mixed together, no reaction will take place without ignition. The explosion that occurs upon ignition is an example of a reaction that proceeds through the generation of unstable, short-lived intermediates called free radicals, molecules with an odd number of electrons. These molecules cannot have closed shells no matter how the electrons are shared, because all electrons must be paired in a closed shell. Free radicals often attack other molecules to grab the missing electron and obtain a closed shell. It is interesting to note that, when a free radical reacts with a non-radical, the reaction products must always include a free radical, since the total number of valence electrons remains odd. Only when two free radicals react with each other can the products all be non-radicals. In the case of 2 H2 + O2, H atoms are generated by the ignition process, and a series of reactions occurs in which the reactants are a free radical and a molecule, in which the net effect of all the reactions is production of H2O. Reactions involving free radicals will be discussed more fully in a later chapter. Reactions involving gases, or gases and liquids, are most important in the chemistry of the atmosphere. The Law of mass action states that the rate for a reaction depends on the concentrations of the reactants. Concentration is a measure of the quantity of reactant per unit volume, for example, the number of moles in a liter of air or water, or the number of molecules in a cubic meter, etc. The rate for the reaction represents the change in concentration in a given time, for example, the number of moles per liter of products generated per second. For a simple reaction where a molecules of A react with b molecules of B to make products, the law of mass action states that rate at which products are generated is Rate = k [A]a [B]b, where the brackets […] denote standard chemical notation for concentration. The parameter k is the rate constant, that may depend on environmental factors (temperature, pressure, etc) but not on [A] or [B]. 6 In general chemical reactions are reversible. For example, the common industrial reaction of carbon monoxide with water, → CO + H2O ← CO2 + H2 , produces a mixture of all four gases in the reactor. The laws of thermodynamics state that there is a unique relationship between concentrations of reactants and products such that the rates of forward and reverse reactions are equal, and there can be no spontaneous changes in the concentrations. This condition is called equilibrium. If we write a general chemical reaction as a molecules of reactant A, plus b molecules of B, etc., reacting to form x molecules of X, y molecules of Y, etc., then the equilibrium conditions is given by [A]a[B]b... . K= [X]x[Y]y... The parameter K is called the equilibrium constant, which typically depends on environmental parameters just like the rate constant k. 6.3. Absorption of radiation by molecules. Light and radiant heat (infrared radiation) propagate through space as waves, called electromagnetic waves because there are an electric field and a magnetic field associated with each wave (the magnetic field is not important for the purposes of this course). Figure 6.4 illustrates the basic properties of electromagnetic waves. All electromagnetic radiation travels at the speed of light, c=3×108 m s-1 in vacuum. If we could take a snapshot of a light wave as it traveled for 1 s, it would be 3x×108 m long, and would look like the sine wave shown in the figure. The distance between two successive crests on the wave is called the wavelength (denoted λ). The frequency (denoted ν) is the number of wave cycles (wavelengths) that pass a reference point per unit time, and since our snapshot shows exactly the number of peaks that passed in one second, ν is also the number of peaks in the picture, i.e. ν =c/λ. Alternatively, 1/ν is the time it takes the wave to travel one wavelength at speed c. Figure 6.4 Schematic diagram of the time-varying electric field due to an electromagnetic wave as it propagates through space. Electromagnetic radiation, although wave-like in nature, is composed of packets of energy called photons. Thus light is both a wave and a particle. For a given electromagnetic wave of wavelength λ the energy associated with each photon is given by 7 E = hc/λ = hν where h is Planck's constant (h=6.626x10-34 J sec). This was one of Planck's great discoveries; it implies that photons with shorter wavelengths are more energetic than photons with longer wavelengths. Matter can emit radiation if its temperature is above 0 K (absolute zero), and it will emit radiation at the same wavelengths of light that it can absorb. An object that absorbs radiation at all wavelengths incident on it necessarily emits radiation at all wavelengths. This ideal material is called a black body; most solid objects, such as the solid body of the earth, or ocean water, behave almost as black bodies. Planck showed that the intensity of light that is emitted from a black body as a function of wavelength (λ) or frequency (ν), is given by the following function (now called the Planck function): FLUX (λ) = 2πhc2 /λ5 ., hc exp( )-1 λkT where FLUX (units: Watts m-2 m-1; 1 W ≡ 1 J s-1) is the amount of energy in light with wavelengths between λ and λ+∆λ passing through surface with area 1 m2 each second. Planck’s Law indicates that the temperature of an object determines the intensity of radiation emitted by the object at any wavelength, provided that the object can absorb radiation at that wavelength. Planck’s Law is a consequence of the fact that matter and radiation must come into equilibrium if they are enclosed under steady conditions long enough. Figure 6.5 shows the Planck function for objects at several different temperatures. The human eye can see light with wavelengths between about 0.4 to 0.8 micrometers (µm). The figure shows that objects at a temperature of 300 K (about room temperature) emit virtually no photons that are visible. However, objects at 300K or at 200K do emit photons, but wavelengths are far into the infrared, which humans cannot see. When we see an object such as a table, for example, we are seeing visible light that is reflected from the table and that reaches our eyes. The visible light may have originated on the sun, which has a surface temperature of about 5500 K, or from a light bulb with a temperature above 1200 K. These hot objects emit a significant amount of energy in the visible region of the spectrum. 8 Figure 6.5. The Planck function for several temperatures is plotted versus wavelength λ (upper scale, 1 µm = 10-6 m) or wave number (≡ 1/λ = ν/c; wave numbers are proportional to photon energy, like ν, but in units more convenient than frequency). The Planck function gives the energy flux from an object divided up according to wavelength (or frequency), for a given temperature. Long before Planck, however, scientists had determined by direct experiment that the total energy flux from an object, at all wavelengths, depended only on temperature, and they derived an empirical equation called the Stefan-Boltzmann law to describe this relationship: TOTAL ENERGY FLUX = σ T4 . Here the total energy flux (units: W m-2) is shown to vary as the 4th power of the absolute temperature, T (K), with a constant of proportionality σ = 5.67 × 10-8 W m-2 K-4, the Stefan-Boltzmann constant. The Stefan-Boltzmann law was obtained in the 19th century by observing the rate at which real objects lost energy via radiation, with many decades passing before Planck showed that it could be derived from his radiation law. It is possible to do simple demonstrations that allow us to visualize the phenomena described by the Stefan-Boltzmann Law and Planck's equation, even without sophisticated measuring devices. Here is a list of several: 9 • • Take two objects of different materials (e.g. a brick and a steel ball) that look different in reflected light, and placed them in an oven that can reach about 900 C (1200 Kelvin). At this temperature they emit light at a high rate at wavelengths that we can see visually. Even though they looked different in reflected visible light, both objects look the same as they glow under these conditions. In fact it will often be difficult to see them at all inside the oven, which is also glowing. This experiment illustrates that ordinary solid objects emit and absorb radiation more or less like "black" bodies, which is the same as saying that their emission spectra follow Planck's equation. Put a strong prism in front of a lantern slide projector to disperse the light, and varied the temperature of the lamp by changing the applied voltage (for example, using a variable transformer). The rapid disappearance of the blue light will be apparent as the temperature is lowered (and vice versa), as will changes in the total amount of light coming from the projector. This experiment allows us to visualize Planck’s function directly and illustrates the phenomena that Planck sought to explain. The changes in emission rate at various wavelengths relate directly to our understanding of sunlight and of heat radiation from the earth. 10 Main points of Chapter 6 1. Atoms form chemical bonds by rearranging electrons in the outer (valence) shell to localize the electrons between the nuclei. 2. Elements form chemical compounds in a reproducible pattern captured in the Periodic Table. 3. The most stable compounds are those in which all electrons are paired, and each atom in the molecule attains a closed shell configuration, either by sharing pairs with its neighbors or by exchanging electrons. The arrangements of electrons in shells, and the rules for sharing or exchanging between neighboring atoms in a molecule, lie behind the pattern observed in the Periodic Table. 4. Rates for chemical reactions follow the Law of Mass Action. 5. There is a unique relationship between concentrations of reactants and products that correspond to equilibrium. All chemical systems tend to evolve towards equilibrium and concentrations cannot change spontaneously once equilibrium is reached. 6. Light may be regarded as both a propagating electric field in the shape of a sine wave and as particles called photons. The relationship between the speed of light (c), its wavelength (λ), and its frequency (ν), is c = λν. 7. Every photon has a specific energy proportional to its frequency, or inversely proportional to its wavelength, E = hc/λ = hν. 8. Matter emits radiation depending on its temperature. The total flux of radiation emitted is given by the Stefan-Boltzmann equation, Flux (W m-2) = σT4, where σ is the Stefan-Boltzmann constant, 5.67x10-8 W m-2 K-4. The flux as a function of 2πhc2 /λ5 (W m-2 m-1). wavelength is given by FLUX (λ) = hc exp( )-1 λkT 9. Matter can emit light only at wavelengths that it can absorb. 10. Most atmospheric gases can neither emit nor absorb light at the long wavelengths (infrared) emitted by cold objects, such as the Earth. Those relatively rare atmospheric molecules that can absorb infrared radiation have asymmetric distribution of charge (e.g. a dipole, like the water molecule) that causes the molecules to experience a force due to the oscillating electric field of the light. 11