Download Chapter 6: The basics of chemistry and interaction of

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