MODULE 3
THE CHEMISTRY OF THE ENVIRONMENT
At the end of the module, you should be able to:
TLO 7: Apply chemistry to the environment for calculations and design appropriate for
water and air.
TLO 8: Apply chemistry to the environment for calculations and design appropriate for soil
Introduction
Environmental chemistry is a branch of chemistry that deals with the study of effect of
chemicals on environment. These include the formation of compounds, how chemicals get
into the environment, the changes they undergo once introduced into the environment,
the number of chemicals in the environment and how they enter the organisms and other
things from the environment and the damage they cause.
In other words, environmental chemistry deals with the effect of pollutants on the
environment and the ways and means to reduce the contamination. This branch of
chemistry is the study of pollutants behavior from the environmental point of view. The
environmental chemistry is further classified into main areas; measurement of pollutants,
and study of the behavior of the contaminants.
Environmental chemistry is vital as chemicals introduced into the environment proves to be
harmful not only to the environment but also to human health and economy. We have
discussed below why environmental chemistry is essential and how it benefits human
health, environment and the economy.
Unit 1
The Chemistry of Water
UNIT LEARNING OUTCOMES
TLO 7: Apply chemistry to the environment for calculations and design appropriate for
water and air.
ENGAGE
In its purest form, it's odorless, nearly colorless and tasteless. It's in your body, the food you
eat and the beverages you drink. You use it to clean yourself, your clothes, your dishes,
your car and everything else around you. You can travel on it or jump in it to cool off on
hot summer days. Many of the products that you use every day contain it or were
manufactured using it. All forms of life need it, and if they don't get enough of it, they die.
Political disputes have centered around it. In some places, it's treasured and incredibly
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difficult to get. In others, it's incredibly easy to get and then squandered. What substance is
more necessary to our existence than any other? Water.
At its most basic, water is a molecule with one oxygen atom and two hydrogen atoms,
bonded together by shared electrons. It is a V-shaped polar molecule, which means that
it's charged positively near the hydrogen atoms and negatively near the oxygen atom.
Water molecules are naturally attracted and stick to each other because of this polarity,
forming a hydrogen bond. This hydrogen bond is the reason behind many of water's
special properties, such as the fact that it's denser in its liquid state than in its solid state (ice
floats on water). We'll look closer at these special properties later.
Water is the only substance that occurs naturally as a solid (ice), a liquid and a gas (water
vapor). It covers about 70 percent of the Earth for a total of approximately 332.5 million
cubic miles (1,386 million cubic kilometers) [source: U.S. Geological Survey]. If you're familiar
with the lines "Water, water, everywhere, nor any drop to drink" from the poem "The Rime of
the Ancient Mariner," you'll understand that most of this water -- 97 percent of it -- is
undrinkable because it's saltwater. Only 3 percent of the world's water supply is freshwater,
and 77 percent of that is frozen. Of the 23 percent that is not frozen, only a half a percent is
available to supply every plant, animal and person on Earth with all the water they need to
survive [source: National Geographic].
So water is pretty simple, right? Actually, there are a lot of things about it that scientists still
don't fully understand. And the problem of making sure that enough clean, drinkable
water is available to everyone and everything that needs it is anything but simple. In this
article, we'll look at some of these problems. We'll also explore exactly what plants, animals
and people do with water and learn more about what makes water so special.
EXPLAIN
WATER
A chemist's view of the world is not as
narrow as one might think! Yes, we start with
the atom, and then go on to the rules
governing the kinds of structural units that
can be made from them. We are taught
early on to predict the properties of bulk
matter from these geometric arrangements.
And then we come to H2O, and are shocked to find that many of these predictions are
way off, and that water (and by implication, life itself) should not even exist on our planet!
But we soon learn that this tiny combination of three nuclei and eight electrons possesses
special properties that make it unique among the more than 15 million chemical species
we presently know. When we stop to ponder the consequences of this, chemistry moves
from being an arcane science to a voyage of wonder and pleasure as we learn to relate
the microscopic world of the atom to the greater world in which we all live.
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The molecule of water
A molecule is an aggregation of atomic nuclei and
electrons that is sufficiently stable to possess observable
properties — and there are few molecules that are more
stable and difficult to decompose than H2O. In water,
each hydrogen nucleus is bound to the central oxygen
atom by a pair of electrons that are shared between
them; chemists call this shared electron pair a covalent
chemical bond.
In H2O, only two of the six outer-shell electrons of oxygen are used for this purpose, leaving
four electrons which are organized into two non-bonding pairs. The four electron pairs
surrounding the oxygen tend to arrange themselves as far from each other as possible in
order to minimize repulsions between these clouds of negative charge. This would ordinarily
result in a tetrahedral geometry in which the angle between electron pairs (and therefore
the H-O-H bond angle) is 109.5°. However, because the two non-bonding pairs remain
closer to the oxygen atom, these exert a stronger repulsion against the two covalent
bonding pairs, effectively pushing the two hydrogen atoms closer together. The result is a
distorted tetrahedral arrangement in which the H—O—H angle is 104.5°.
Although the water molecule carries no net electric
charge, its eight electrons are not distributed uniformly;
there is slightly more negative charge (purple) at the
oxygen end of the molecule, and a compensating
positive charge (green) at the hydrogen end. The
resulting polarity is largely responsible for water's unique
properties.
Because molecules are smaller than light waves, they cannot be observed directly, and
must be "visualized" by alternative means. This computer-generated image comes from
calculations that model the electron distribution in the H2O molecule. The outer envelope
shows the effective "surface" of the molecule as defined by the extent of the cloud of
negative electric charge created by the eight electrons.
Hydrogen bonding
The H2O molecule is electrically neutral, but the positive
and negative charges are not distributed uniformly. This is
illustrated by the gradation in color in the schematic
diagram here. The electronic (negative) charge is
concentrated at the oxygen end of the molecule, owing
partly to the nonbonding electrons (solid blue circles),
and to oxygen's high nuclear charge which exerts
stronger attractions on the electrons. This charge
displacement constitutes an electric dipole, represented
by the arrow at the bottom; you can think of this dipole
as the electrical "image" of a water molecule.
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As we all learned in school, opposite charges attract, so
the partially-positive hydrogen atom on one water
molecule is electrostatically attracted to the partiallynegative oxygen on a neighboring molecule. This
process is called (somewhat misleadingly) hydrogen
bonding. Notice that the hydrogen bond (shown by the
dashed green line) is somewhat longer than the
covalent O—H bond. This means that it is considerably
weaker; it is so weak, in fact,that a given hydrogen bond
cannot survive for more than a tiny fraction of a second.
Sticky, Wet Water
Water has some unusual properties due to its hydrogen bonds. One property is cohesion,
the tendency for water molecules to stick together. The cohesive forces between water
molecules are responsible for the phenomenon known as surface tension. The molecules at
the surface do not have other like molecules on all sides of them and consequently they
cohere more strongly to those directly associated with them on the surface. For example, if
you drop a tiny amount of water onto a very smooth surface, the water molecules will stick
together and form a droplet, rather than spread out over the surface. The same thing
happens when water slowly drips from a leaky faucet. The water doesn't fall from the
faucet as individual water molecules but as droplets of water.
Another important physical property of water, is adhesion. In terms of water, adhesion is
the bonding of a water molecule to another substance, such as the sides of a leaf's veins.
This process happens because hydrogen bonds are special in that they break and reform
with great frequency. This constant rearranging of hydrogen bonds allows a percentage of
all the molecules in a given sample to bond to another substance. This grip-like
characteristic that water molecules form causes capillary action, the ability of a liquid to
flow against gravity in a narrow space. An example of capillary action is when you place a
straw into a glass of water. The water seems to climb up the straw before you even place
your mouth on the straw. The water has created hydrogen bonds with the surface of the
straw, causing the water to adhere to the sides of the straw. As the hydrogen bonds keep
interchanging with the straw's surface, the water molecules interchange positions and
some begin to ascend the straw.
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Adhesion and capillary action are necessary to the survival of most organisms. It is the
mechanism that is responsible for water transport in plants through roots and stems, and in
animals through small blood vessels.
Hydrogen bonds also explain why water’s boiling point (100°C) is higher than the boiling
points of similar substances without hydrogen bonds. Because of water’s relatively high
boiling point, most water exists in a liquid state on Earth. Liquid water is needed by all living
organisms. Therefore, the availability of liquid water enables life to survive over much of the
planet.
Furthermore, water has a high specific heat because it takes a lot of energy to raise or
lower the temperature of water. As a result, water plays a very important role in
temperature regulation. Since cells are made up of water, this property helps to maintain
homeostasis.
The melting point of water is 0°C. Below this temperature, water is a solid (ice). Unlike most
chemical substances, water in a solid state has a lower density than water in a liquid state.
This is because water expands when it freezes. Again, hydrogen bonding is the reason.
Hydrogen bonds cause water molecules to line up less efficiently in ice than in liquid water.
As a result, water molecules are spaced farther apart in ice, giving ice a lower density than
liquid water. A substance with lower density floats on a substance with higher density. This
explains why ice floats on liquid water, whereas many other solids sink to the bottom of
liquid water.
In a large body of water, such as a lake or the ocean, the water with the greatest density
always sinks to the bottom. Water is most dense at about 4°C. As a result, the water at the
bottom of a lake or the ocean usually has temperature of about 4°C. In climates with cold
winters, this layer of 4°C water insulates the bottom of a lake from freezing temperatures.
Lake organisms such as fish can survive the winter by staying in this cold, but unfrozen,
water at the bottom of the lake.
SOLUTIONS
Most of the materials that we encounter in everyday life are mixtures. Many mixtures are
homogeneous; that is, their components are uniformly intermingled on a molecular level.
Homogeneous mixtures are called solutions. Examples of solutions abound in the world
around is. The air we breathe is a solution of several gases. Brass is solid solution of zinc and
copper. The fluids that run through our bodies are solutions, carrying a great variety of
essential nutrients, salts, and other materials.
Solutions are extremely important. throughout the living world, solutions are necessary for
maintenance and survival. In the human body, nutrients are transported in solution, while
waste products are removed as solutions. In plants, all internal organs are constantly
bathed in moisture; in photosynthesis, a wet surface is required on which gases can diffuse
materials to the body.
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A solution, since it is homogeneous, will have the same properties all throughout No matter
where we take a sample from a solution, we will obtain materials with same composition as
that of any sample from the solution.
SOLUTION - is a homogeneous mixture of two or more substances
* Homogeneous - only one phase; no part is separated by another by a detectable
boundary.
2 components
1. solute - dissolved substance; present in smaller amount
2. solvent - dissolving medium; present in greater amount
Aquation or Hydration - when water is used as a solvent
SOLUBILITY - amount of solute that can be dissolved in a given amount of solvent at a
specific temperature; usually expressed as grams of solute per 100 grams of solvent
soluble - substance that dissolves in the solvent
insoluble - substance that does not dissolve in the solvent
solute solubility (grams solute/100 grams solvent)
less than 0.1
0.10 - 1.0
1.0 - 10
greater than 10
qualitative solubility description
insoluble
slightly soluble
soluble
very soluble
I. FACTORS AFFECTING SOLUBILITY
A. SOLID in LIQUID
1. Shaking or Agitation
2. Powdering or Pulverizing
3. Temperature
4. Nature of Reactants - “like dissolves like”
B. LIQUID in LIQUID
1. Miscible - dissolve in any amount in each other
ex.
alcohol + water; oil+ gasoline/CCl 4
2. Partially miscible - have limited solubility in each other
ex.
tincture of iodine + water
3. Immiscible - do not dissolve in each other and forms two separate layers upon mixing
ex.
oil + water
C. GAS in LIQUID
1. Pressure - affects gases only
* Henry’s Law - the weight of a gas dissolved by a given amount of solvent is directly
proportional to the pressure exerted by the as wen in equilibrium with the solution.
2. Temperature
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II. TYPES OF SOLUTIONS
A. According to state or phase
1. Gaseous Solution
Solute
Solvent
gas
gas
liquid
gas
solid
gas
2. Liquid Solution
Solute
Solvent
gas
liquid
liquid
liquid
solid
liquid
3. Solid Solution
Solute
gas
liquid
solid
Solvent
solid
solid
solid
Example
air
water vapor in air
dust particle in air
Example
carbonated drinks
alcohol in water
sugar in water
Example
hydrogen in palladium metal
amalgam
alloy
B. According to solubility
1. Unsaturated solution - a solution that can still take more of the solute in addition to what
it already contains; less concentrated than a saturated solution
2. Saturated Solution - contains the maximum amount of solute that can be dissolved in the
given amount of solvent; contains the maximum amount of dissolved solute
necessary for the existence of an equilibrium between dissolved and undissolved
solute
3. Supersaturated solution - contains more of the solute that it normally contains at an
elevated temperature; more concentrated than a saturated solution
C. According to the amount of dissolved solute
1. Concentrated solution - contains greater amount of solute
2. Dilute solution - contains lesser amount of solute; concentrated solution + water
III. WHY SUBSATANCES DISSOLVE? SOLUTION FORMATION
A. solute - solute attraction - attraction between solute particles
B. solvent - solvent attraction - attraction between solvent particles
* driving force for solution formation
solute - solvent attraction - attraction between solute and solvent particles
IV. METHODS OF EXPRESSING CONCENTRATION OF SOLUTIONS
concentration - amount of solute present in a specified amount of solvent or solution
A. PERCENTAGE METHODS
1. Percentage by weight or mass (%w/w)
- most frequently used by chemists
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(
(
)
)
2. Percentage by volume (%v/v)
- used when both solute and solvent are liquids
3. Percentage by weight-volume (%w/v)
- usually encountered in hospitals and industrial settings
- used when working with a solid solute and liquid solvent
B. MOLE FRACTION (yi)
C. MOLARITY (M) or FORMALITY (F)
- most often used in chemical laboratories
M = number of gram-molecular weight (mol) of solute per liter of solution
F = number of gram-formula weight (mol) of solute per liter of solution
D. NORMALITY (N)
- number of gram-equivalent weights of solute per liter of solution
- most often encountered in neutralization reactions
f = factor (equivalent/mol)
1. acid: f - number of replaceable or ionizable H+
ex.
HCl
f=1
H2SO4
f=2
H3PO4
f=3
HC2H3O2
f=1
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2. base: f - number of replaceable or ionizale OHex.
NaOH
f=1
Ca(OH)2
f=2
Al(OH)3
f=3
3. salt: f - total + or - valence
ex.
NaCl
f=1
CaBr2
f=2
Al2S3
f=6
E. MOLALITY (m)
- used in experimental situations where changes in temperature are concerned
- number of moles of solute per kilogram (kg) of solvent
V. DILUTION CONCEPT
A common problem encountered when working with solutions in the laboratory is
that of diluting a solution of known concentration (usually called a stock solution) to a
lower concentration.
Dilution - a process in which more solvent is added to a solution in order to lower the
concentration of the solution. The amount of solute present is now distributed in a
larger amount of solvent.
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EXPLAIN
Activity 1:
1.
2.
3.
The Unique Properties of Water
Self-Assessment No. 1
Enumerate and discuss the unique properties of water. Give examples for each.
Water is the so-called universal solvent. Give everyday examples of these aqueous
solutions.
How would you able to increase the degree of saturation of a sugar solution?
To be submitted in Google classroom on:
ELABORATE & EVALUATE
Activity 2:
Solutions
Self-Assessment No. 2
1. Compute the weight (g) of solute needed to make the solutions listed below:
a. 1250 mL of 0.85 M AlCl3
b. 0.500 L of 9.25 N K2SO4
c. 350 g of 13.2 %w/w of Ca(OH)2
2. Calculate the %w/v of a solution made by dissolving 22.0 g CH 3OH (methanol) in
C2H5OH (ethanol) to make 100 mL solution.
3. What is the %v/v of 10.0 mL of CH 3OH (methanol) dissolved in water to a volume of
40.0 mL?
4. A solution is prepared at 25oC by mixing 20 g of K2SO4 with 150 cc of water. The density
of water at this temperature is 1 g/cc. What is the m of K2SO4?
5. Express the formality (F) the concentration of each of the following solutions:
a. 28.4% NH3 by weight, having a density of 0.808 g/mL
b. 69.5% HNO3 by weigt, aving a density of 1.42 g/mL
6. An aqueous solution of hydrofluoric acid is 12.5 molal with respect to HF and has a
density of 1.070 g/mL at 25oC. Calculate
a. mole fraction of HF
b. %w/w of HF
c. molarity
d. molality
7. Calculate the volume of concentrated reagent required to prepare the diluted
solutions indicated:
a. 15 M NH3 to prepare 50 mL of 6.0 M NH3
b. 18 M H2SO4 to prepare 250 mL of 10.0 M H2SO4
8. Calculate the molarity of the solutions by mixing 250 mL of 0.75 M H 2SO4 with
a. 150 mL of H2O
b. 250 mL of 0.70 M H2SO4
To be submitted in Google classroom on:
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