This document provides an overview of gases and the gas laws. It begins with an introduction to the composition of Earth's atmosphere and how gases shield us and keep the planet warm. It then discusses gas pressure and how Torricelli invented the barometer to measure atmospheric pressure in 1643. Other sections cover the gas laws of Boyle, Charles, Avogadro, and the ideal gas law. The document provides examples of using these laws to solve gas stoichiometry problems. It concludes with an introduction to Dalton's law of partial pressures and the kinetic molecular theory of gases.
What constitutes waste depends on the eye of the beholder; one person's waste can be a resource for another person.[1] Though waste is a physical object, its generation is a physical and psychological process.[1] The definitions used by various agencies are as below.
United Nations Environment Program
According to the Basel Convention on the Control of Transboundary Movements of Hazardous Wastes and Their Disposal of 1989, Art. 2(1), "'Wastes' are substance or objects, which are disposed of or are intended to be disposed of or are required to be disposed of by the provisions of national law".[2]
United Nations Statistics Division
The UNSD Glossary of Environment Statistics[3] describes waste as "materials that are not prime products (that is, products produced for the market) for which the generator has no further use in terms of his/her own purposes of production, transformation or consumption, and of which he/she wants to dispose. Wastes may be generated during the extraction of raw materials, the processing of raw materials into intermediate and final products, the consumption of final products, and other human activities. Residuals recycled or reused at the place of generation are excluded."
European Union
Under the Waste Framework Directive 2008/98/EC, Art. 3(1), the European Union defines waste as "an object the holder discards, intends to discard or is required to discard."[4] For a more structural description of the Waste Directive, see the European Commission's summary.
Types of Waste
Municipal Waste
The Organization for Economic Co-operation and Development also known as OECD defines municipal solid waste (MSW) as “waste collected and treated by or for municipalities”. [5] Typically this type of waste includes household waste, commercial waste, and demolition or construction waste. In 2018, the Environmental Protection Agency concluded that 292.4 tons of municipal waste was generated which equated to about 4.9 pounds per day per person. Out of the 292.4 tons, approximately 69 million tons were recycled, and 25 million tons were composted. [6]
Household Waste and Commercial Waste
Household waste more commonly known as trash or garbage are items that are typically thrown away daily from ordinary households. Items often included in this category include product packaging, yard waste, clothing, food scraps, appliance, paints, and batteries.[7] Most of the items that are collected by municipalities end up in landfills across the world. In the United States, it is estimated that 11.3 million tons of textile waste is generated. On an individual level, it is estimated that the average American throws away 81.5 pounds of clothes each year.[8] As online shopping becomes more prevalent, items such as cardboard, bubble wrap, shipping envelopes are ending up in landfills across the United States. The EPA has estimated that approximately 10.1 million tons of plastic containers and packaging ended up landfills in 2018. The EPA noted that only 30.
Unit 8 - Information and Communication Technology (Paper I).pdfThiyagu K
This slides describes the basic concepts of ICT, basics of Email, Emerging Technology and Digital Initiatives in Education. This presentations aligns with the UGC Paper I syllabus.
The Indian economy is classified into different sectors to simplify the analysis and understanding of economic activities. For Class 10, it's essential to grasp the sectors of the Indian economy, understand their characteristics, and recognize their importance. This guide will provide detailed notes on the Sectors of the Indian Economy Class 10, using specific long-tail keywords to enhance comprehension.
For more information, visit-www.vavaclasses.com
Palestine last event orientationfvgnh .pptxRaedMohamed3
An EFL lesson about the current events in Palestine. It is intended to be for intermediate students who wish to increase their listening skills through a short lesson in power point.
How to Make a Field invisible in Odoo 17Celine George
It is possible to hide or invisible some fields in odoo. Commonly using “invisible” attribute in the field definition to invisible the fields. This slide will show how to make a field invisible in odoo 17.
Operation “Blue Star” is the only event in the history of Independent India where the state went into war with its own people. Even after about 40 years it is not clear if it was culmination of states anger over people of the region, a political game of power or start of dictatorial chapter in the democratic setup.
The people of Punjab felt alienated from main stream due to denial of their just demands during a long democratic struggle since independence. As it happen all over the word, it led to militant struggle with great loss of lives of military, police and civilian personnel. Killing of Indira Gandhi and massacre of innocent Sikhs in Delhi and other India cities was also associated with this movement.
This is a presentation by Dada Robert in a Your Skill Boost masterclass organised by the Excellence Foundation for South Sudan (EFSS) on Saturday, the 25th and Sunday, the 26th of May 2024.
He discussed the concept of quality improvement, emphasizing its applicability to various aspects of life, including personal, project, and program improvements. He defined quality as doing the right thing at the right time in the right way to achieve the best possible results and discussed the concept of the "gap" between what we know and what we do, and how this gap represents the areas we need to improve. He explained the scientific approach to quality improvement, which involves systematic performance analysis, testing and learning, and implementing change ideas. He also highlighted the importance of client focus and a team approach to quality improvement.
Students, digital devices and success - Andreas Schleicher - 27 May 2024..pptxEduSkills OECD
Andreas Schleicher presents at the OECD webinar ‘Digital devices in schools: detrimental distraction or secret to success?’ on 27 May 2024. The presentation was based on findings from PISA 2022 results and the webinar helped launch the PISA in Focus ‘Managing screen time: How to protect and equip students against distraction’ https://www.oecd-ilibrary.org/education/managing-screen-time_7c225af4-en and the OECD Education Policy Perspective ‘Students, digital devices and success’ can be found here - https://oe.cd/il/5yV
How to Split Bills in the Odoo 17 POS ModuleCeline George
Bills have a main role in point of sale procedure. It will help to track sales, handling payments and giving receipts to customers. Bill splitting also has an important role in POS. For example, If some friends come together for dinner and if they want to divide the bill then it is possible by POS bill splitting. This slide will show how to split bills in odoo 17 POS.
3. Intro
Earth’s atmosphere is a gaseous
solution that consists mainly of
nitrogen (N2) and oxygen (O2). This
atmosphere supports life and acts as
a waste receptacle for many
industrial processes. The chemical
reactions that follow often lead to
various types of pollution, including
smog and acid rain.
4. Intro
The gases in the atmosphere also shield
us from harmful radiation from the sun
and keep the earth warm by reflecting
heat radiation back toward the earth. In
fact, there is now great concern that an
increase in atmospheric carbon dioxide,
a product of the combustion of fossil
fuels , is causing a dangerous warming
of the earth.
6. Pressure
Gas uniformly fills a container, is
easily compressed, and mixes
completely with any other gas.
One of the most important
properties is that it exerts pressure
on its surroundings equally.
11. Pressure
Barometer - A device to
measure atmospheric pressure,
was invented in 1643 by Torricelli
(a student of Galileo).
Torricelli’s barometer is
constructed by filling a glass
tube with liquid mercury and
inverting it in a dish of mercury.
12. Pressure
Barometer - A device to
measure atmospheric pressure,
was invented in 1643 by Torricelli
(a student of Galileo).
Notice that a large quantity of
mercury stays in the tube. In
fact, at sea level the height of
this column of mercury averages
760 mm.
13. Pressure
Barometer - A device to
measure atmospheric pressure,
was invented in 1643 by Torricelli
(a student of Galileo).
Atmospheric pressure results
from the mass of the air being
pulled toward the center of the
earth by gravity.
14. Pressure
Barometer - A device to
measure atmospheric pressure,
was invented in 1643 by Torricelli
(a student of Galileo).
Atmospheric pressure varies
with weather changes and
altitude.
15. Pressure
Manometer – an instrument for measuring pressure
often below that of atmospheric pressure.
16. Units of Pressure
Pressure = force/area
torr – in honor of Torricelli is equal
to a mm Hg.
760 mm Hg = 760 torr
1 atm = 760 torr
Pascal = N/m2
1atm = 101,325 Pa
19. Boyles Law
Boyle (1627-1691) –
performed the first quantitative
experiments on gases. Used a
J tube to measure pressures.
PV = k; P1V1=P2V2
k is a constant for a given
sample of air at a specific
temperature.
20. Boyles Law
Pressure and volume are often plotted.
P vs V – gives a hyperbola and an
inverse relationship.
Boyles law rearranged is
V=k/P=k1/P; when plotted as V vs
1/P – gives a straight line with the
intercept of zero
22. Boyles Law
Boyles’s law holds precisely at
very low temperatures, but
varies at higher pressures. PV
will vary as pressure is varied.
An ideal gas is a gas that
strictly obeys Boyles’ law.
24. Charles Law
Charles (1746-1823) – the first person to fill a
balloon with hydrogen gas and who made
the first solo balloon flight.
Charles found in 1787 that the volume of a
gas at constant pressure increases linearly
with the temperature of the gas.
V = bT ; V1/T1 = V2/T2
T is in Kelvin, b is a proportionality constant
25. Charles Law
Temperature vs Volume plots a
straight line.
Slope will vary will type of gas.
All gas plots of T vs V will
extrapolate to zero at the same
temperature.
-273° or 0 K
27. Avogadro’s Law
Avogadro (1811) – postulated
that equal volumes of gases at
the same temperature and
pressure contain the same
number of particles (moles).
V = an; V1/n1=V2/n2
n is number of moles; a is a
proportionality constant.
29. Ideal Gas Law
The relationships that Boyle, Charles
and Avogadro presented can be
combined to show how the volume of a
gas depends on pressure, temperature,
and number of moles of gas present.
V = R(Tn/P)
R is the universal gas constant.
A combination of proportionality
constants
30. Ideal Gas Law
The equation is often rearranged to
form the more common:
PV=nRT
R=.08206 Latm/Kmol
31. Ideal Gas Law
Limitations
A gas that obeys this equation is said
to behave ideally. The ideal gas
equation is best regarded as a limiting
law, it expresses behavior that real
gases approach at low pressures and
high temperatures.
Most gases behave ideally at pressures
below 1 atm.
32. Example Problem
A sample of hydrogen gas (H2) has a
volume of 8.56 L at a temperature of
0°C and a pressure of 1.5 atm.
Calculate the moles of H2 molecules
present in this gas sample.
33. Example Problem
V = 8.56 L at a temperature of 0°C
P =1.5 atm
Calculate the moles
PV=nRT; R = .08206 L atm/K mol
.57 mol
34. Example Problem 2
You have a sample of ammonia gas with
a a volume of 7.0ml at a pressure of
1.68 atm. The gas is compressed to a
volume of 2.7 ml at a constant
temperature. Use the ideal gas law to
calculate the final pressure.
35. Example Problem 2
V1 = 7.0 ml at a pressure of 1.68 atm
V2 =2.7 ml at a constant temperature.
Calculate the final pressure.
PV=nRT; but nRT are constant: PV=PV
4.4 atm
36. Example Problem 3
A sample of methane gas that has a
volume of 3.8 L at 5°C is heated to 86°C
at constant pressure. Calculate its new
volume.
37. Example Problem 3
V1 = 3.8 L and T15°C
T2=86°C at constant pressure.
Calculate its new volume.
PV=nRT but n, R and P are constant: V1/T1 = V2/T2
4.9 L
38. Example Problem 4
A sample of diborane gas (B2H6), a
substance that burst into flame when
exposed to air, has a pressure of 345 torr
at a temperature of -15°C and a volume
of 3.48 L. If conditions are changed so
that the temperature is 36°C and the
pressure is 468 torr, what will be the
volume of the sample.
39. Example Problem 4
P1=345 torr, T1=-15°C and V1 = 3.48L
T2=36°C and P2=468 torr,
What is the volume?
PV = nRT; nR are constant: P1V1/T1 =
P2V2/T2
3.1 L
40. Example Problem 5
A sample containing 0.35 mol argon gas at
a temperature of 13°C and a pressure of
568 torr is heated to 56°C and a pressure
of 897 torr. Calculate the change in
volume that occurs.
41. Example Problem 5
n=0.35, T1=13°C, P1=568 torr
T2=56°C, P2=897 torr
Calculate the change in volume that occurs.
-3 L
PV=nRT; R = .08206 L atm/K mol
45. Gas Stoichiometry
We use STP or standard
temperature and pressure of an ideal
gas to make calculations with a gas.
1 atm
0°C (273K)
1 mole = 22.42 L becomes a
conversion factor for dimensional
analysis.
46. Gas Stoichiometry
Example
A sample of nitrogen gas has a volume
of 1.75 L at STP. How many moles of
N2 are present?
1.75L N2 x
1mole N2
22.42L N2
=
7.81 x 10-2
mol N2
47. Gas Stoichiometry
Example 2
Quicklime (CaO) is produced by the
thermal decomposition of calcium
carbonate (CaCO3). Calculate the
volume of CO2 at STP produced from
the decomposition of 152g CaCO3 by
the reaction
CaCO3(s) CaO(s) + CO2(g)
48. Gas Stoichiometry
Example 2
152g CaCO3
22.42 L = 1 mol of gas at STP
Calculate the volume of CO2
CaCO3(s) CaO(s) + CO2(g)
34.1 L CO2 at STP
49. Gas Stoichiometry
Example 3
A sample of methane gas having a
volume of 2.80 L at 25°C and 1.65
atm was mixed with a sample of
oxygen gas having a volume of 35.0 L
at 31°C and 1.25 atm. The mixture
was then ignited to form carbon
dioxide and water. Calculate the
volume of CO2 formed at a pressure of
2.50 atm and a temperature of 125°.
50. Gas Stoichiometry
Example 3
CH4 V=2.80 L at 25°C and 1.65 atm
Oxygen V=35.0 L at 31°C and 1.25 atm.
Calculate the volume of CO2 at 2.50 atm and
125°C.
The mixture was then ignited to form carbon
dioxide and water.
51. Gas Stoichiometry
Example 3
CH4 V=2.80 L at 25°C and 1.65 atm
Oxygen V=35.0 L at 31°C and 1.25 atm.
Calculate the volume of CO2 at 2.50 atm and
125°C.
CH4(g) + O2(g) CO2(g) + H2O(g)
2.47 L
53. Molar Mass of a Gas
One use of the ideal gas law is in the calculations of the
molar mass of a gas from its measured density.
n =
P
D
P =
grams of gas
molar mass
=
m
molar mass
=
nRT
V
=
m / molar mass( ) RT
V
=
mRT
V(molar mass)
=
m
V
dRT
molar mass
; Molar mass =
dRT
P
54. Gas Density/Molar
Mass Example
The density of a gas was measured
at 1.50 atm and 27°C and found to
be 1.95 g/L. Calculate the molar
mass of the gas.
32.0 g/mol
56. John Dalton
John Dalton formed his atomic theory from his
experiments and studies of the mixture of gases.
His observations car be summarized as follows:
For a mixture of gases in a container, the total
pressure exerted is the sum of the pressures
that each as would exert if it were alone.
57. John Dalton
Ptotal=P1+P2+P3+….
Subscripts refer to the individual gases
and Px refers to partial pressure that a
particular gas would exert if it were
alone in the container.
Each Partial pressure can be derived
from the ideal gas law and added
together to determine the total.
58. John Dalton
Ptotal=P1+P2+P3+….
Since each partial pressure can be
broken down into ; the Ptotal can be
represented by:
Ptotal=
Ptotal=
nx RT
V
(n1 + n2 + n3 + ...)
RT
V
÷
ntotal
RT
V
÷
59. For a mixture of ideal gases, it is the total number of
moles of particles that is important, not the identity
or composition of the involved gas particle.
60. Dalton’s Law Example
Mixtures of helium and oxygen can be
used in scuba diving tanks to help
prevent “the bends.” For a particular
dive, 46 L He at 25° and 1.0 atm and
12 L O2 at 25° and 1.0 atm were
pumped into a tank with a volume of
5.0 L. Calculate the partial pressure of
each gas and the total pressure in the
tank at 25° C.
61. Mole Fraction
The ratio of the number of
moles of a given component
in a mixture to the total
number of moles in the
mixture.
= nχ x/ntotal
= Pχ 1/Ptotal
62. Dalton’s Law Example
The partial pressure of oxygen was
observed to be 156 torr in air with a
total atmospheric pressure of 743
torr. Calculate the mole fraction of O2
present.
63. Dalton’s Law Example
The mole fraction of nitrogen in the air
is 0.7808. Calculate the partial
pressure of N2 in air when the
atmospheric pressure is 760 torr.
64. Collecting Gas over
Water
A mixture of gases results whenever a
gas is collected by displacement of water.
In this situation, the gas in the bottle is a
mixture of water vapor and the oxygen
being collected.
65. Collecting Gas over
Water
Water vapor is present because
molecules of water escape from the
surface of the liquid and collect in the
space above the liquid.
66. Collecting Gas over
Water
Molecules of water also return to the
liquid. When the rate of escape equals
the rate of return, the number of water
molecules in the vapor state remain
constant.
67. Collecting Gas over
Water
When the number of water molecules in
the vapor state remain constant the
pressure of the water vapor remains
constant.
69. Collecting Gas over
Water Example
A sample of solid potassium chlorate
(KClO3) was heated in a test tube and
decomposed by the reaction:
2KClO3(s) → 2KCl(s) + 3O2(g)
70. Collecting Gas over
Water Example
The oxygen produced was collected by
displacement of water at 22°C at a total
pressure of 754 torr. The volume of gas
collected was .650L, and the vapor
pressure of water at 22°C is 21 torr.
Calculate the partial pressure of O2 in the
gas collected and the mass of KClO3 in
the sample that was decomposed.
2KClO3(s) → 2KCl(s) + 3O2(g)
71. Collecting Gas over
Water Example
oxygen T=22°C and V=.650L
Total pressure = 754 torr.
vapor pressure at 22°C is 21 torr.
Calculate the partial pressure of O2 and
the mass of KClO3 in the sample
2KClO3(s) → 2KCl(s) + 3O2(g)
2.59 x 10-2
mol O2 2.12 g KClO3
73. Kinetic Molecular Theory
KMT
A simple model that attempts to
explain the properties of an ideal
gas. This model is based on
speculations about the behavior of
the individual gas particles (atoms or
molecules).
74. Kinetic Molecular Theory
KMT
1. The particles are so small
compared with the distances
between them that the volume of
the individual particles can be
assumed to be negligible (zero).
75. Kinetic Molecular Theory
KMT
2. The particles are in constant
motion. The collisions of the
particles with the walls of the
container are the cause of the
pressure exerted by the gas.
76. Kinetic Molecular Theory
KMT
3. The particles are assumed to exert
no forces on each other; they are
assumed neither to attract nor to
repel each other.
77. Kinetic Molecular Theory
KMT
4. The average kinetic energy of a
collection of gas particles is
assumed to be directly
proportional to the Kelvin
temperature of the gas.
78. KMT and Boyle’s Law
Because a decrease in volume, the
gas particles will hit the walls more
often, thus increasing the pressure
79. KMT and Charles Law
When the gas is heated to a higher
temperature, the speeds of its molecules
increase and thus hit the walls more
often and with more force. Volume
and/or pressure will increase.
80. KMT and Advogadro’s Law
An increase in the number of gas
particles at the same temperature would
cause the pressure to increase if the
volume were constant.
81. KMT and Advogadro’s Law
The volume of a gas (at constant T and
P) depends only on the number of gas
particles present. The individual particles
are not a factor because the particle
volumes are so small compared with the
distances between the particles.
82. KMT and Dalton’s Law
All gas particles are independent of each
other and that the volumes of the
individual particles are unimportant.
Identities of the gas particles do not
matter.
83. The Meaning of Temperature
Kelvin temperature indicates the
average kinetic energy of the gas
particles.
The exact relationship between
temperature and average kinetic
energy can be expressed:
(KE)avg=3/2 RT
84. The Meaning of Temperature
The Kelvin temperature is an index
of the random motions of the
particles of a gas, with higher
temperature meaning greater
motion.
85. Root Mean Square Velocity
u2
=the average of the squares of
the particle velocities.
The square root of u2
is called the
root mean square velocity and is
symbolized with urms
urms= =
M= mole of gas particles (kg)
R = ; J = kgm2
/s2
3RT
Mu2
8.3145
J
Kimol
86. Root Mean Square Velocity
Example
Calculate the root mean square
velocity for the atoms in a sample
of helium gas at 25°C.
1.36 x 103
m/s
87. Mean Free Path
The average distance a particle travels
between collisions in a particular gas
sample.
1 x 10-7
m for O2 at STP
urms=500 m/s
88. Mean Free Path
A velocity distribution that show the effect
of temperature on the velocity distribution
in a gas.
91. Effusion
Effusion describes the
transfer of gas from one
chamber to another (usually
through a small hole or
porous opening).
The rate of transfer is said
to be the rate of effusion.
92. Effusion
The rate of effusion of a gas is
inversely proportional to the square
root of the mass of its particles.
93. Effusion
Rate of Effusion for gas 1
Rate of Effusion for gas 2
=
M2
M1
Temperature must be the same for
both gases.
M represents the molar masses of the
gases.
Units can be in g or kg since the
units will cancel out.
This is called Graham’s law of effusion:
94. Effusion Example
Calculate the effusion rates of
hydrogen gas (H2) and Uranium
hexafluoride (UF6), a gas used in the
enrichment process to produce fuel
for nuclear reactors.
13.2 : 1
96. Real Gases
An ideal gas is a hypothetical
concept. No gas exactly
follows the ideal gas law,
although many gases come
very close at low pressures
and/or high temperatures.
97. Real Gases
Thus ideal gas behavior can
best be thought of as the
behavior approached by real
gases under certain
conditions.
98. Real Gases
Plots of PV/nRT vs. P for several gases
(200K). Ideal behavior only at low
pressures.
99. Real Gases
Plots of PV/nRT vs. P for N2 at three
temperatures. Ideal behavior at higher
temperatures.
100. KMT Modifications
Johannes van der Walls (1837-
1923), a physics professor at the
University of Amsterdam started
work in the area of ideal vs real gas
behavior. He won the nobel prize in
1910 for his work.
101. KMT Modifications
van der Waals modifications to the ideal
gas law accounted for the volume of
particle space. Therefore adjusting for the
volume actually available to a give gas
molecule.
V-nb
n is number of moles
b is an empirical constant
102. KMT Modifications
van der Waals modifications to the ideal
gas law allowed for the attractions that
occur among particle in a real gas which
is dependent upon the concentration of
the particles.
, pressure correction
a is proportionality constant.
Pobs = P'
− a
n
v
÷
2
103. van der Waals Equation
Insert both corrections and the equation
can be written as:
Rearranged for van der Waals:
Pobs =
nRT
V − nb
− a
n
V
÷
2
Pobs + a
n
V
÷
2
x V − nb( ) = nRT
104. van der Waals
Equation
a and b values are
determined for a given gas
by fitting experimental
behavior. That is a and b
are varied until the best fit
of the observed pressure is
obtained under all
conditions.
106. van der Waals
Ideal behavior at low
pressure (large volume)
makes sense because the
small amount of volume that
the particles consume are not
a factor.
107. van der Waals
Ideal behavior at high
temperatures also makes
sense because particles are
moving at such a high rate
that their interparticle
interactions are not very
important.
109. Chemistry in the
Atmosphere
The most important gases to us are
those in the atmosphere that surround
the earth’s surface.
The principal components are N2 and
O2, but many other important gases,
such as H2O and CO2, are also
present.
110.
111. Chemistry in the
Atmosphere
Because of gravitational effects, the
composition of the earth’s atmosphere
is not constant; heavier molecules
tend to be near the earth’s surface,
and light molecules tend to migrate to
higher altitudes, with some eventually
escaping into space.
112. Chemistry in the
Atmosphere
The chemistry occurring in the higher
levels of the atmosphere is mostly
determined by the effects of high-
energy radiation and particles from
the sun and other sources in space.
The upper atmosphere serves as a
shield to prevent this radiation from
reaching earth.
113. Chemistry in the
Atmosphere
The troposphere (closest to earth) is
strongly influenced by human
activities. Millions of tons of gases
and particulates are released into the
troposphere by our highly industrial
civilization.
114. Chemistry in the
Atmosphere
Severe air pollution is found around
many large cities. The two main
sources of pollution are transportation
and the production of electricity. The
combustion of petroleum in vehicles
produces CO, CO2, NO, NO2.
115. Chemistry in the
Atmosphere
The complex chemistry of polluted air
appears to center around the nitrogen
oxides (NOx). At high temperatures
found in the gasoline and diesel
engines of cars and trucks, N2 and O2
react to form a small quantity of NO
that is emitted into the air with the
exhaust gases. NO is immediately
oxidized in air to NO2.
116. Reactions in the
Atomsphere
NO2(g)
radiant
energy
→ NO(g) + O(g)
O(g) + O2(g) → O3(g)
Ozone is very reactive and can react directly with
other pollutants, or the ozone can absorb light and
break up to form an energetically excited O2 molecule
(O2*) and excited O (O*).
117. Reactions in the
Atomsphere
O*
+ H2O → 2OH
OH + NO2 → HNO3
The end product of this whole process is often referred
to as photochemical smog, so called because light is
required to initiate some of the reactions.
118. Reactions in the
Atomsphere
S(coal) + O2(g) → SO2(g)
2SO2(g ) + O2 →2SO3(g )
SO3(g) + H2O(l) → H2SO4(aq)
Sulfuric acid is very corrosive to both living things and
building materials. Another result of this type of
pollution is called acid rain.