Environmental chemistry (notes)12IEEM BY Muhammad Fahad Ansari 12IEEM14

Environmental Chemistry


What is Chemistry?

The science that deals with matter; its composition; its properties; the changes in composition that it will undergo; its relationship to energy; and the laws, principles, theories, and concepts that describe, interpret, and predict its behavior and basic nature.

Periodic Table

Metals: They are good conductors of electricity; they contain only one or two weakly electrons in their outer energy levels.

Non-metals: They have relatively strong attraction for the electrons in their outer energy levels.

Metalloids: The elements whose properties are intermediate between those of the metallic and non-metallic elements.

Bonding between Elements:

Ionic Bond: The electrostatic attraction that exists between positive ions formed when one atom loses electrons and negative ions formed by the atoms that gain the electrons.

Covalent Bond (Coordinate Covalent Bond): A bond between two atoms in which they are overlapping orbitals and sharing a pair of electrons, both of which originally belonged to only one of the atoms.

The Polar Covalent Bond

When, in a covalent bond, certain atoms have a partial positive charge and others have a partial negative charge, we say that the covalent bond is polar. This type of bond is called a polar covalent bond. (One atom “hogs” the electrons; the atom that “hogs” is the one with greater electronegativity).

The Metallic Bond

A metallic bond results when two metals bond. In metallic bonding, the metal atoms donate valence electrons to become cations. These valence electrons are not directly transferred to another atom as they are in ionic bonding. Instead, they move about freely throughout the sample, producing an attractive force that keeps the metal cations in place. Often, the behavior of these electrons is referred to as a “sea of mobile electrons.” Because of the motions of the free electrons, metals are characteristically good conductors of electricity and heat.

• Ionic bond with: exchange of electrons to form stable octet; bond between a metal and nonmetal; example: NaCl, CaF2

• Covalent bond with: sharing of electrons to form stable octet; bond between nonmetals; example: Cl2, N2

• Polar covalent bond with: unequal sharing of electrons to form stable octet; example: H2O, NH3

• Metallic bond with: sea of mobile electrons; bond between metals; example: Cu, Ag

Muhammad Fahad Ansari 12IEEM14

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Molecules can also be polar or Non-polar

• Any diatomic molecule that has a non-polar bond is non-polar. Example: All diatomic molecules: Cl 2, N2, O2, etc… Otherwise, it will generally be polar.

• Molecules that consist of 3 or more atoms are generally polar unless the following condition is met: if the central atom has no lone pairs and is surrounded by atoms of one element, then the molecule will be non-polar, for example, CO2. In these cases, individual bond polarities cancel each other.

Electro-negativity: (Table)

• The bonds between two elements can be purely ionic, purely covalent, or anywhere in between depending on the elements involved.

• The electro-negativity expresses the power of an atom in a molecule to attract shared electrons.

• The electro-negativity values come from measurements of the strength of bonds between atoms and from measurements of the amount of electrical energy required to remove an electron from an atom of an element.

• Fluorine (F) which has the strongest attraction for electrons of all the reactive elements is assigned an electro-negativity value of 4. All other elements are compared to fluorine.

• According to Pauling (1901), if the difference between the electro-negativity values of the two elements is 2.0 or more, the bond is mainly ionic. If the electro-negativity difference is less than 2.0, the bond is mainly covalent.

Organic and inorganic compounds

The distinction between the compounds found in the earth associated with rocks and soil and those associated with living organisms.

The two groups of compounds differ in their occurrence and in many of their properties. These differences provided the basis for one of the earliest efforts to classify compounds.

The compounds associated with living organisms called organic compounds, and the earthy compounds are called inorganic compounds.

Organic chemistry is simply the chemistry of the compounds of the clement carbon. Except the carbon-containing compounds such as carbon dioxide, the carbonates, bicarbonates, and carbides are not considered to be organic compounds.

In other words the properties of organic compounds are largely a result of the covalent bond, and that inorganic compounds, which have different properties, get their characteristics from the ionic bond.

Table- Comparative Properties of Ionic and Covalent Compounds


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Inorganic (Ionic) Compounds Organic (Covalent) Compounds1 Solids, often crystalline,

relatively hard and brittleMay be solid, liquid, or gas, depending on the size of the molecules and the degree of intermolecular attraction (polarity and ability to form hydrogen bonds)

2 Relatively high melting and boiling points

Relatively low melting and boiling points, depending primarily on the size of the molecules and the amount of intermolecular attraction

3 Relatively soluble in water and insoluble in organic (non polar) liquids

Relatively insoluble in water and soluble in organic liquids; polarity and ability to form hydrogen bonds greatly increases water solubility

4 Electrolytes-solutions conduct electricity

Most are non electrolytes

5 Relatively stable to heat Relatively unstable to heat

Biochemistry: The chemistry that deals with the substances and reaction in living organisms.

Electrochemistry: the chemistry dealing with the effects of electricity on chemical reactions and the production of electricity through chemical processes. Chemical reactions involve an electron exchange between the reactants. These reactions are called oxidation-reduction reactions or electrochemical reactions. There are two distinctly different classes of electrochemical reactions: those that produce electrical energy and those that are produced by electrical energy.

Chemical Reactions

When a chemical reaction occurs, at least one product is formed that is different from the substances present before the change occurred. As an example, it is possible to pass an electric current through a sample of water and obtain a mixture of oxygen and hydrogen gases. That change is a chemical reaction because neither oxygen nor hydrogen was present as elements before the change took place.

Any chemical change involves two sets of substances: reactants and products. A reactant is an element or compound present before a chemical change takes place. In the example above, only one reactant was present: water. A product is an element or compound formed as a result of the chemical reaction. In the preceding example, both hydrogen and oxygen are products of the reaction.

Chemical reactions are represented by means of chemical equations. A chemical equation is a symbolic statement that represents the changes that occur during a chemical reaction. The statement consists of the symbols of the elements and the formulas of the products and reactants, along with other symbols that represent certain conditions present in the reaction. For example, the arrow (or yields) sign, *, separates the reactants from the products in a reaction. The chemical equation that represents the electrolysis of water is 2H2O → 2H2 + O2.


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Types of Chemical Reactions

Most chemical reactions can be categorized into one of about five general types: synthesis, decomposition, single replacement, double replacement, and oxidation-reduction. A miscellaneous category is also needed for reactions that do not fit into one of these five categories.

Characteristics of each type

Synthesis: Two substances combine to form one new substance:

In general: A + B → AB

For example:

2 Na + Cl 2 → 2 NaCl or CaO + H 2 O → Ca(OH) 2

Decomposition: One substance breaks down to form two new substances:

In general: AB → A + B

For example: 2 H 2 O → 2 H 2 + O 2

Single Replacement: An element and a compound react such that the element replaces one other element in the compound:

In general: A + BC → AC + B

For example: Mg + 2 HCl → MgCl 2 + H 2

Double Replacement: Two compounds react with each other in such a way that they exchange partners with each other:

In general: AB + CD → AD + CB

For example:

NaBr + HCl → NaCl + HBr

Oxidation-reduction: One or more elements in the reaction changes its oxidation state during the reaction: In general: A 3+ → A 6+

For example: Cr 3+ → Cr 6+

Energy changes and chemical kinetics

Chemical reactions are typically accompanied by energy changes. The equation for the synthesis of ammonia from its elements is N 2 + 3 H 2 → 2 NH 3 , but that reaction takes place only under very special conditions—namely at a high temperature and pressure and in the presence of a catalyst. Energy changes that occur during chemical reactions are the subject of a field of science known as thermodynamics.

In addition, chemical reactions are often a good deal more complex than a chemical equation might lead one to believe. For example, one can write the equation for the synthesis of hydrogen iodide from its elements, as follows: H 2 + I 2 → 2 HI. In fact, chemists know that this reaction does not take place in a single step. Instead, it occurs in a series of reactions in which


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hydrogen and iodine atoms react with each other one at a time. The final equation, H 2 + I 2 → 2 HI, is actually no more than a summary of the net result of all those reactions. The field of chemistry that deals with the details of chemical reactions is known as chemical kinetics.

Reaction of sulfuric acid and sugar. The acid dehydrates the sugar forming a pillar of carbon (black) and steam.

Reaction Rates and Chemical Equilibrium

There are few questions related to factors that influence reaction rates

• Why must substances such as gasoline be heated before they will start burning?

• To what does the term catalytic refer in the catalytic conversion process used in auto emission control systems?

• Why do wood shavings burn more rapidly than pieces of wood?

The answers to these and other questions involve the rate of chemical reactions.

Forward and Reverse Reaction:

Ammonia (NH3) manufactured as a fertilizer

N2 + 3H2 → 2NH3

2NH3 → N2 + 3H2

N2 + 3H2 ↔ 2NH3

Reactants Products

When the rate of the forward reaction equals the rate of the reverse reaction, the system is said to be in chemical equilibrium.

Chemical Equilibrium: Exists when two opposing reactions are occurring simultaneously and at the same rate. The maintenance of body weight is an example of a kind of equilibrium. Reaction rates and equilibrium are interrelated.

Factors Affecting the Rate of Chemical Reactions

All the factors involved in the transformation of reactants to products in chemical changes are called chemical kinetics. The major factors that affect the rate of reactions are the nature of the reactants, temperature, catalysts, and the reactants concentration.


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Nature of the reactants: Different substances react at different rates. In general, the readiness with which any two substances react together depends largely on the structure of the atoms of the elements and the nature of the bonds that hold the atoms together in the substance.

Temperature: In general, an increase in temperature results in an increase in rate of reaction. Before reaction can occur between two substances, the bonds already existing between the atoms of each substance must be weakened and broken. An increase in temperature results in greater kinetic energy, and thus motion, of the individual atoms of each substance. The greater energy of motion causes the atoms to tend to pull away from one another. As a result, the existing bonds are weakened or broken. The energy that must be added to weakened bonds so that substances will react is called the energy of activation.

Many reactions give off energy, the energy liberated by such reactions is usually in the form of heat. Some times radiant energy (for example, light in combustion reactions) and electrical energy may also be given off. Reactions that give off (evolve) heat are called exothermic reactions. All combustion reactions are exothermic reactions such as:

• C + O2 → CO2 + heat

• Mg + F2 → MgF2 + heat

• N2 + 3H2 → 2NH3 + heat

Reactions that absorb heat are called endothermic reactions. The following equations are examples of endothermic reactions:

• 2KClO3 + heat → 2KCl + 3O2

• CaCO3 + heat → CaO + CO2

• 2NH3 + heat → N2 + 3H2

One reaction of equilibrium will evolve heat (exothermic) and the other reaction will absorb heat (endothermic).

ExothermicN2 + 3H2 -------------------→ 2NH3 + heat

←------------------ Endothermic

Catalysts:

A catalyst is a substance that changes the rate of a reaction without itself being used up in the process. Catalysts function in several different ways, but their overall effect is to change the reaction rate by lowering the energy of reactants activation.


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Reactants Concentration:

In general, increasing the concentration of reactants increases the rate of reaction. For example, substances burn more rapidly in pure oxygen than in air, which is only about 20 percent oxygen.

The rate of the reaction is proportional to the product of the concentration of the reactants. But the rate of the reaction is equal to the product of the concentration of the reactants times a proportionality constant (k). The constant K takes into account the effect of the nature of the reactants, temperature, pressure, and catalyst on the reaction.

Rate of Chemical Reaction

The rate or speed with which reactants disappear or a product appears.

The rate at which the concentration of one of the reactants decreases, or one of the products increases with time.

The decomposition of dinitrogen pentoxide (N2O5) in an inert solvent carbon tetrachloride (CCl4)

A typical unit for a rate of reaction (mol per L per s)

2 N2O5 (in CCl4) → 2N2O4 (in CCl4) + O2 (g)

One product, O2, is a gas that is virtually insoluble in the reaction mixture.

Decomposition of N2O5 (in CCl4) at 45 0C [N2O5] = 1.40 MTime, s Total volume O2, cm3

(at STP)0 0432 1.32753 2.181116 2.891582 3.631986 4.102343 4.46. .. .. .∞a 5.93

Reaction Kinetics: Reaction Mechanisms


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Terms Activation Energy - The difference in energy between the reactants and the transition state that is the

energy barrier the reactants must overcome to achieve a chemical reaction.

Catalyst - A substance that lowers the activation energy for a chemical reaction without being

chemically altered by the reaction.

Elementary Step - A reaction that represents a single collision or intramolecular step in a reaction

mechanism.

Homogeneous Catalyst - A catalyst that is in the same phase as the reactants.

Intermediate - A species that is both produced and consumed in a chemical reaction. As such, it does

not appear in the overall reaction but is proposed to be produced in one elementary step and consumed

in another.

Kinetics - The study of the rate and mechanism of chemical reactions.

Mechanism - The series of elementary steps that combine to produce the path molecules take from

reactant(s) to product(s) in a chemical reaction.

Order - In the rate law of a reaction, the power to which the concentration of a reagent is raised. Or,

the sum of the powers on the concentration terms in the rate law.

Rate - The speed of a reaction measured in amount or reagent consumed or product produced per unit

time.

Rate Constant - The proportionality constant in the rate law expression. This factor is a measure of the

intrinsic reactivity of the reaction but is not constant with respect to temperature.

Rate Law - An expression of the dependence of the rate of a reaction on the concentrations of

reactants.

Rate Limiting Step - The slowest elementary step in a mechanism. The rate of the reaction must equal

the rate of the slowest step because the reaction can go no faster than its slowest step.

Reaction Coordinate Diagram - A plot of free energy versus the reaction coordinate for a reaction that

provides a pictorial representation of the lowest energy path from reactants to products.

Steady-State Approximation - The assumption that the rate of formation and consumption of a highly

reactive intermediate are equal so that the change in intermediate concentration with respect to time is

approximated to be zero.

Transition State - The species with the highest energy between reactants and products on a reaction

coordinate diagram, it is a short-lived species that represents a combination of product-like and

reactant-like properties.

Chemical Mechanisms

By describing how atoms and molecules interact to generate products, mechanisms help us to

understand how the world around us functions at a fundamental level. A mechanism is a series of


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elementary steps whose sum is the overall reaction. An elementary step is a reaction that is meant to

represent a single collision or vibration that leads to a chemical change. For a mechanism to be

considered valid, its sum must equal the overall balanced equation, its predicted rate law must agree

with experimental data, and its predictions of intermediates must not be contrary to experimental

observations. A mechanism may never be proven because we cannot ever see a chemical reaction--both

the time scale of an elementary step and the size of atoms are too small. Furthermore, we must guess at

the identity of many intermediates because they are usually so reactive that they can not be isolated.

Instead, a chemist proposes reaction mechanisms and tests their validity against experimental data,

ruling out mechanisms that are inconsistent with results. These experiments may be strategically

designed to trap an intermediate product to prove its existence as a stepping-point in the total reaction.

To aid in our understanding of mechanisms, we will draw reaction coordinate diagrams that trace the

free energy path of a reaction from reactants to products. The activation energy of a reaction represents

the difference in energy between the reactants and the highest point on a reaction coordinate diagram.

We will derive the Arrhenius Equation, which relates the rate constant for a reaction to its activation

energy. Local minima on the reaction coordinate diagram are positions occupied by intermediates. By

comparing the reaction coordinate diagram for a catalyzed and a uncatalyzed process, we can see that

catalysts function by altering the route the reaction takes from reactants to products without the

catalyst being altered.

Properties of Mechanisms

Mechanisms describe in a stepwise manner the exact collisions and events that are required for the

conversion of reactants into products. Mechanisms achieve that goal by breaking up the overall

balanced chemical equation into a series of elementary steps. An elementary step is written to mean a

single collision or molecular vibration that results in a chemical reaction. The following picture of an

elementary step shows a single collision between water and boron trifluoride:


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Figure %: Schematic representation of an elementary step

The molecularity of an elementary step describes the number of reactive partners in the elementary

step. For example, the above elementary step is called bimolecular because two molecules collide.

Commonly, elementary steps are mono-, bi-, or ter-molecular. The probability of four molecules

colliding at exactly the same place and time is so small that we can safely assume that no reaction will

ever be tetra-molecular. Because take up a large amount of space, we will represent elementary steps in

this Spark Note as normal reactions with molecular formula line equations. You will know from the

context (i.e. talking about the steps of a mechanism) whether the reaction is an elementary step or an

overall reaction.

To better understand mechanisms, let's consider the following mechanism for the decomposition of

ozone, O3:

The above mechanism exhibits a property of all mechanisms: it is a series of elementary steps whose

sum is the overall balanced reaction. Note the presence of the oxygen atom, O, intermediate in the


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above equation. It is an intermediate because it is both created and destroyed in the mechanism and

does not appear in the net equation.

Another property of mechanisms is that they must predict the experimentally determined rate law. To

calculate the rate law from a mechanism you need to first know the rate limiting step. The rate limiting

step determines the rate of the reaction because it is the slowest step. You can rationalize that a

reaction can only go so fast as its slowest step by thinking about what happens when you encounter an

accident on the highway that closes all but one lane. You may have been able to race along at 65 m.p.h.

(depending on your state's laws) before you reached the lane closure but the slow passage of cars past

the accident limits your rate. You can only go as fast through that one lane as the slowest car in front of

you.

In the above, the first reaction is labeled as "slow". This reaction is the rate determining step because it

is the slowest step. As we have stated, that means that the rate of the overall reaction is equal to the

rate of the rate determining step. The rate of an elementary step is the rate constant for that step

multiplied by the concentrations of the reactants raised to their stoichiometric powers. Note that this

rule only applies for elementary steps. The rate of an overall reaction is NOT the product of the

concentrations of the reactants raised to their stoichiometric powers. The rate law for the first

elementary step in the is rate = k [O3]. Because this step is the rate determining step, the rate law is also

the rate law for the overall reaction. Using similar techniques we can calculate the rate law predicted by

any mechanism. We then check the predicted rate law against the experimentally determined rate law

to test the validity of the proposed mechanism.

Reaction Coordinate Diagrams

We can follow the progress of a reaction on its way from reactants to products by graphing the energy

of the species versus the reaction coordinate. We will be vague in describing the reaction coordinate

because its definition is a mess of other variables composed to best make sense of the progress of the

reaction. The value of the reaction coordinate is between zero and one. Understanding the meaning of

the reaction coordinate is not important, just know that small values of reaction coordinate (0-0.2)

mean little reaction has taken place and large values (0.8-1.0) mean that the reaction is almost over. It is

a kind of scale of the progress of a reaction. A typical reaction coordinate diagram for a mechanism with

a single step is shown below:


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Figure %: A reaction coordinate diagram for a single-step reaction

Note that the reactants are placed on the left and the products on the right. The choice of the energy

levels of the reactants and products is dictated by their energies, those with higher energies are higher

on the diagram and those with lower energies are lower on the diagram. The difference is energy

between the reactants and the transition state is called the activation energy. The activation energy is

the height of the energy barrier of the reaction. The transition state is the point of maximum energy on

the diagram which represents a species possessing both reactant-like and product-like properties.

Because it is so high in energy, the transition state is very reactive and can never be isolated due to its

extremely short lifetime. The relative energy of the reactants and products, the ΔE on the diagram,

determines whether the reaction is exothermic or endothermic. A reaction will be exothermic if the

energy of the products is less than the energy of the reactants. A reaction is endothermic when the

energy of the products is greater than the energy of the reactants. The is for an exothermic reaction.

Below is a reaction coordinate diagram for an endothermic reaction.

Figure %: Reaction coordinate diagram for an endothermic reaction

If a reaction has n elementary steps in its mechanism, there will be n–1 minima between the products

and reactants representing intermediates. There will also be n maxima representing the n transition

states. For example, a reaction with three elementary steps could have the following reaction

coordinate diagram.


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Figure %: Reaction coordinate diagram for a three-step reaction

One confusing point about reaction coordinate diagrams is how to determine what the rate determining

step is. Even experienced chemists consistently get this type of problem wrong. The rate determining

step is not the one with the highest activation energy for the step. The rate determining step is the step

whose transition state has the highest energy.

Activation Energy and the Arrhenius Equation

Intuitively, it makes sense that a reaction with a higher activation barrier will be slower. Think of how

much harder you must roll a ball up a large hill than a smaller one. Let's consider chemical reactions

more deeply to derive an equation which describes the relationship between the rate constant of a

reaction and its activation barrier. To simplify our derivation, we will assume that the reaction has a one-

step mechanism. This elementary step represents a collision as shown in . Therefore, the frequency of

the collisions, f, will be important in our equation. Notice that only a certain orientation of the molecules

will lead to a reaction. For example, the following collision will not lead to a reaction. The reagent

molecules simply bounce off of one another:


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Figure %: Only specific orientations during a collision will lead to a reaction.

Therefore, we will need to include an orientation factor (or steric factor), p, that takes into account the

fact that only a certain fraction of collisions will lead to reaction due to the orientation of the molecules.

Another factor we must consider is that only a certain fraction of the molecules colliding will have

enough energy to overcome the activation barrier. The Boltzmann distribution is a thermodynamic

equation that tells us what fraction of the molecules have a certain amount of energy. As you know, at

higher temperatures the average kinetic energy of the molecules increases. Therefore, at higher

temperatures more molecules have energy greater than the activation energy--as shown:

Figure %: Boltzmann distributions for T1 greater than T2

Combining the above considerations, we state the following relationship between the rate constant and

the activation energy, called the Arrhenius equation:


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The variable k is the rate constant, which is dependent on the frequency of the collisions f, orientation

factor p, activation energy Ea, and temperature T. From the expression for the Arrhenius equation you

should note that a small increase in activation energy leads to a large decrease in rate constant.

Furthermore, temperature has a similarly exponential effect on the rate constant. An experimental rule

of thumb is that a 10oC increase in temperature leads to a doubling of the rate constant.

One application of the Arrhenius equation that is useful is the determination of the activation energy for

a reaction. Taking the natural log of the Arrhenius equation gives a linear equation:

A graph of ln k versus 1 / T should give a straight line whose slope is - Ea / R. By measuring the rate

constant at a range of different temperatures, you can construct a graph to determine the activation

energy of a reaction.

Catalysis

A catalyst speeds up a reaction without being explicit in the overall balanced equation. It does this by

providing an alternate mechanism for the reaction that has a lower activation barrier than does the

uncatalyzed pathway. Compare the catalytic and regular mechanisms for the hydrogenation of ethylene

to ethane and their associated reaction coordinate diagrams in :


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Figure %: Mechanisms of ethylene hydrogenation

As you can see, the catalyst changes the mechanism of the reaction and lowers the activation energy.

The catalyst, because it does not appear in the overall balanced equation has absolutely no effect on the

thermodynamics of the reaction.

There are two types of catalysts--heterogeneous catalysts and homogeneous catalysts. There is no

fundamental difference in how these catalysts work. The difference lies in whether the catalyst is in the

same phase (solid, liquid, or gas) as the reagents. A homogeneous catalyst is in the same phase as the

reactants while a heterogeneous catalyst is not. An enzyme is a biological homogeneous catalyst.

Acid Base Reactions

Inorganic Compounds

Centuries ago chemists began to classify substances by grouping them into elements, compounds, and mixtures.

Then the elements were classifies further into metals and nonmetals. Likewise, compounds were divided into organic and inorganic.

The organic compounds were eventually grouped into smaller classes such as alcohol, hydrocarbons, ketones, ethers, and so on.

The inorganic compounds were further classified as acids, bases, slats and so on.


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Classes of inorganic Compounds:

Compounds were classified on the basis of their properties. However, theories can help you understand why the compounds of a given class have similar properties, such as atomic structure and the bonding.

Acids-Bases-Buffers

Importance of acids, Bases and Buffers (salts)

It is fact that, stomach contains acids. Overeating can cause a person to suffer from frequent bouts of acid indigestion. So that you can always have something on hand to neutralize this inevitable excess stomach acid (Asprin is an acid (acetyl salicylic acid) which may upset the stomach unless properly buffered. Acids, bases, and buffers are important to the functioning of the human body. Many foods we eat either contain acids or yield acids as they are digested.

Yet the acid base balance of your body blood must be maintained with very narrow limits, or you will die. The concentration of H+ (acidity) of the blood should be approximately 4X10-8 M.

If it increases to approximately 6X10-8 M

Or decreases to approximately 2.5X10-8 M death results

So that blood acidity is maintained at the proper level.

Acids, bases and buffers are also of tremendous importance to industry and commerce.

Such as HCl, HNO3, and H2SO4 are manufactured in million-ton quantities annually for a wide variety of uses.

Theories about Acids:

• Santé Arrhenius (1859-1927) Swedish chemist explains why water solutions of certain substances conduct electricity.

Arrhenius proposed the theory of ionization.

• He suggested that these substances, which are called electrolytes, dissociate in water to produce charged particles (ions). The ions are responsible for conducting electricity so that

• Acids are electrolytes when they react with metals to yield hydrogen gas (H2).

Arrhenius defined acids as substances that ionize in water to yield hydrogen ions (H+)

The thing common to all acids is their ability to yield hydrogen ions (H+) in water.

For example: HCl in H2O Yields H+ + Cl-

HF in H2O H+ + F-

HBr in H2O H+ + Br-


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• Bronsted-Lowry Concept about Acids:

According to this theory, acids are defined simply as proton donors.

The hydrogen ion is a proton: the hydrogen atom minus its electron.

So acid-base reactions are not limited to reactions between water solutions. Any substance which gives up a proton (hydrogen ion H+) in any reaction is in an acid. This includes all the substances those are acids according to the Arrhenius concept and many others.

Acids Composition:

In terms of composition, acids consist of the nonmetallic element hydrogen covalently bonded to other nonmetallic element.

The system of naming the acids centers around the number of oxygen atoms present.

The most common oxygen containing acids of the nonmetal has three oxygen atoms include halogens (Cl, Br, I) nitrogen (N), and carbon (C).

These acids are HClO3 chloric acid, HBrO3 bromic acid, HIO3 iodic acid, HNO3 nitric acid and H2CO3

carbonic acid.

Acids of sulfur (S) and phosphorous (P) contain four oxygen atoms H2SO4 sulfuric acid and H3PO4

phosphoric acid.

Only sulfur and phosphorous contain four oxygen atoms the rest contain three oxygen atoms.

The number of hydrogen atoms is the same in all the acids of a given nonmetal.

Acids of phosphorous contain three hydrogen atoms. The acids of sulfur and carbon contain two hydrogen atoms, all the rest contain only one hydrogen atom.

Properties of Acids:

The acids occur in all three physical states (gases, liquids, and solids) HCl as a gas, Acetic acid as a liquid and Boric acid as a solid; in the pure liquid form they do not conduct electricity.

• Acids have a sour taste. For example, the sour taste of lemons and other citrus fruits is to due to citric acid, one of the organic acids.

• Acids react with active metals to yield hydrogen. Sodium (Na), potassium (K), calcium (Ca), zinc (Zn), and other active metals react with acids such as hydrochloric and sulfuric to yield hydrogen gas.

• Acids react with bases to yield a salt and water. Bases and salts are two other major classes of inorganic compounds. Hydrochloric acid (HCl) reacts with the base sodium hydroxide (NaOH) to yield ordinary table salt, which is sodium chloride (NaCl), and water.

• In their pure state some acids are gases, some are liquids, and others are solids.


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• Acids are electrolytes; they conduct electricity when dissolved in water. However, in the liquid state pure acid compounds such as HCl do not conduct electricity.

• Acids turn blue litmus red. Litmus is an organic dye that is used as an indicator. It becomes red when exposed to acids and blue when exposed to bases, so it is a convenient way to distinguish between acids and bases.

Properties of Bases:

Arrhenius and Bronsted-Lowry concepts: Arrhenius considered bases to be substances that ionize in water to yield hydroxide ions (OH-)

The common bases are metallic hydroxides such as NaOH sodium hydroxide, calcium hydroxide Ca(OH-)2. The metallic hydroxide bases are solids that will conduct electricity in molten state-ionically bonded.

NaOH in water Na+ + OH-

Ca (OH-)2 in water Ca2+ + 2OH-

NH3 + H2O NH4OH NH4+ + OH- (NH3 is a nonmetallic hydroxide)

Bronsted-Lowry theory defines bases as protons (H+) acceptors.

Water acts as a base in the presence of an acid. It accepts a proton (H+) from HCl and becomes a hydronium ion (H3O+). Hydronium ion can give up the proton (H+) and the chloride ion (Cl-) can accept a proton, they are also an acid and base, respectively.

NH3 + H2O NH4+ + OH-

Composition of Bases:

The bases are named by simply naming the metal and adding the word hydroxide, such as sodium hydroxide (NaOH), calcium hydroxide Ca (OH-)2, and aluminum hydroxide Al (OH-)3, are the more common bases of the non-transition metals.

Bronsted-Lowry Concept: the solvent enters into the reaction, and acids and bases occur in solution are called conjugate pairs. The hydronium ion is the conjugate acid of the base water, and the chloride ion is the conjugate base of the acid HCl.

Many transition metals may lose varying number of electrons as they react to form compounds. Metals may have more then one oxidation number. For example, copper (Cu) may lose one or two electrons, forming two different ions: Cu1+, Cu2+. [Cu OH and Cu (OH-)2] are possible.

The ions of metallic element of higher charge (oxidation number) indicate the –ic suffix as cupric, and –ious indicate the lower charge (oxidation number).

Salts:

Much of the earth’s crust consists of salts. Acids react with bases to form salts and water.


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HCl + NaOH NaCl + H2O

Composition and properties of salts:

In general, most inorganic salts consist of a metallic element ionically bonded to the nonmetallic element or one of the oxygen containing ions (SO4

2-, CO32-, NO3

- and so on) of the nonmetallic element.

The properties of the salts are essentially the same as those of the ionically bonded compounds.

The salts are crystalline solids that have high melting and boiling points, and they are electrolytes. They are generally more soluble in water and more stable to heat than the organic compounds.

In general binary salts are more stable to heat than the ternary salts. The N-to-H bonds in the ammonium ions are covalent, as are the bonds between oxygen and other nonmetals. Therefore, ammonium salts and salts containing binary anions (negative ions) are less stable to heat than the binary salts.

Oxides:

Oxygen is a special element when combined with hydrogen; it forms the very important compound water (hydrogen oxide). If one hydrogen of water is displaced by a metallic element, the hydroxide (OH -) bases result. Oxygen also combines with both the metallic and nonmetallic elements to form other binary (two-element) compounds. Many of these compounds are salts (compounds of metal and nonmetal), however, because of oxygen special role they are often classified as oxides.

Metallic Oxides:

Oxygen readily combines with the metallic elements to form compounds, most of which are ionically bonded.

Properties of Metallic Oxides:

The metallic oxides have the same general properties as other ionically bonded compounds. They are solids with high melting and boiling points and high heat stability.

Na2O + H2O 2NaOH

CaO + H2O Ca (OH-)2

Nonmetallic Oxides:

Oxygen bonds covalently with the other nonmetallic elements to form oxides. A given nonmetallic element may form more than one oxide, partly because oxygen may bond to the element by both covalent and co-ordinate covalent bonds. If there is more than one atom of nonmetal present, this is


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indicated by prefix (mono, di, tri and so on). System of naming oxides of the nonmetallic elements is as following:

CO carbon monoxide CO2 carbon dioxide

SiO2 silicon dioxide SO2 sulfur dioxide

SO3 sulfur trioxide N2O3 dinitrogen trioxide

P2O5 diphosphrous Pentoxide

Properties of nonmetallic Oxides:

The nonmetallic oxides are covalently bonded compounds, related to organic compounds. Several of the nonmetallic oxides dissolve in water and undergo a reaction with it. The nonmetallic oxides may be considered to be anhydrides of acids. They form acid in water, he soluble nonmetallic oxides are electrolytes.

CO2 + H2O = H2CO3 in (H2O) = H3O+ + HCO3-

SO2 + H2O = H2SO3 in (H2O) = H3O+ + HSO3-

Acid-Base Reactions

Na+OH- + H+NO3- Na+NO3

- + H2O

(aq) (aq) (aq)

Base acid salt water

NH3 + HCl NH4+Cl-

(gas) (gas) (solid) base acid salt

HCl (in water) H+ + Cl-

H2O + HCl H3O+ + Cl-

base acid acid base

The hydronium ion (H3O+) is capable of giving up the proton (to the base) and becoming a water molecule again, it is an acid. Thus the hydronium ion is the conjugate acid of the base water.

Na+OH- (in water) Na+ + OH-

Substances other than metallic hydroxides that are bases when dissolved in water


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NH3 + H2O NH4+ + OH-

Base acid acid base

Acids of halogens

The acids of chlorine

HClO4 Perchloric acid

HClO3 Chloric acid

HClO2 Chlorous acid

HClO hypochlorous acid

HCl Hydrochloric acid

HI Hydroiodic acid

HF Hydrofluoric acid

HBr Hydrobromic acid

Acids of Nitrogen

HNO3 Nitric acid

HNO2 Nitrous acid

Acids of Carbon

H2CO3 Carbonic acid

Acids of Sulfur

H2SO4 Sulfuric acid

H2SO3 Sulfurous acid

Acids of Phosphorous


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H3PO4 Phosphoric acid

H3PO3 Phosphorous acid

Organic acid (weak acid)

HC2H3O2 Acetic acid

Bases of non-transition elements

Li, Na, K (OH-)

Ba, Ca, Mg (OH-)2

Al (OH-)3

Bases of transition elements

Cu1+OH-

Cu2+(OH-)2

Fe2+ (OH-)2

Fe3+ (OH-)3

Breaking Bonds:

Oxidation-Reduction Reactions

Charge transfer reactions - Oxidation/Reduction (Redox) Reactions

Chemical reactions involve making and breaking bonds. So far we have considered making and

breaking ionic bonds between ions in aqueous salt solutions. Now we want to discuss making

and breaking covalent bonds. When we break covalent bonds, there are three ways we can


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separate the two atoms involved in the bond. Consider for example a C-X bond. If when the

bond is broken, the electrons are left with C, C takes on a negative charge and is called a carbon

anion. If the electrons leave with X, then C is left with a positive charge and is a carbon cation.

If one electron stays with each atom the C has no charge, but does not have an octet. The

carbon is called a free radical.

Oxidation-Reduction Reactions:

Chemical reactions involves an electron exchange between the reactants, in which one substance loses electrons and another substance gains electrons or gains a greater share of them these reactions are called oxidation-reduction reactions.

Oxidizing agent: Substance that causes another substance to lose electrons

Reducing agent: Substance that loses electrons in a chemical reaction; it reduces another substance by losing electrons to it.

Chemical Reactions:

Example:Reactants Product

2Na + Cl2 → 2 Na+ Cl-

Loses Gains

Oxidized Reduced

Reducing agent Oxidizing agent

Oxidation Number or Oxidation State:

Reflects the number of electrons lost or gained in relation to the elemental state.


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Simple and arbitrary rules to assign an oxidation number to a substance:

• The oxidation number of an element in its free or uncombined form is zero.

For example: Na0, Mg0, S0, O20, and so on.

• The oxidation number of a mono-atomic (one-atom) cation (positive ion) or anion (negative ion) is equal to its charge. In other words, the oxidation number of mono-atomic ions equals the number of electrons it has lost or gained.

For example: Na1+, Mg2+, Al3+, Cl1-, S2-, and so on.

• The oxidation number of oxygen in compounds is usually -2.

The exceptions are the peroxides, such as H2O2, and compounds of oxygen and fluorine, such as OF2.

• The oxidation number of hydrogen in compounds is usually +1.

The hydrides, such as NaH, are exceptions.

• In the formula for a compound, the sum of the positive oxidation numbers must equal the sum of the negative oxidation numbers.

For example: Mg2+S2-, Na1+Cl1-, K1+Mn7+O48-, H2

2+S6+O48-, and so on.

• In complex ions such as SO42-, PO4

3-, ClO31-, the algebraic sum of the oxidation numbers of the

individual atoms in the ion equals the charge on the ion.

For example: in SO42-, since oxygen is always -2 there will be a total of -8 oxidation number due

to oxygen, therefore, S must have an oxidation number of +6 in order for the ion to have oxidation number of -2.

SO42- x +4(-2) = -2 (net charge on ion)

x -8 = -2

x = 8 -2 = 6

So the oxidation number of S in SO42- , Cl in ClO3

1- and P in PO43- are


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ClO31- x + 3(-2) = -1 (net charge on ion)

X - 6 = - 1

X = 6 – 1 = 5

PO43- x + 4(-2) = -3 (net charge on ion)

X - 8 = - 3 X= 8 – 3 = 5

Oxidation/Reduction (Redox) Reactions

We have just discussed acid/base reactions, which involve proton transfer. Now we will discuss

another kind of charge transfer, electron transfer or oxidation/reduction reactions. In

oxidation/reduction reactions, there is a transfer of charge - an electron - from one species to

another. Oxidation is the loss of electrons and reduction is a gain in electrons. Use these

acronyms to help you remember: Leo Ger - Loss of electrons oxidation; Gain of electrons

reduction or Oil Rig - Oxidation involves loss, Reduction involves gain - of electrons.) These

reactions always occur in pair. That is, an oxidation is always coupled to a reduction. When

something gets oxidized, another agent gains those electrons, acting as the oxidizing agent, and

gets reduced in the process. When a substance gets reduced, it gains electrons from something

that gave them up, the reducing agent, which in the process gets oxidized. It's just like an acid

base reaction. An acid reacts with a base to form a new acid and base.

Reactions in which a pure metal reacts with a substance to form a salt are clearly oxidation

reactions. Consider for example the reaction of sodium metal and chlorine gas.

2Na(s) + Cl2 2NaCl(s)

Na is a pure metal. (Although it really exists as sodium ions surrounded by a sea of electrons),

consider it for our purposes to exist as elemental Na, which has a formal charge of 0. Likewise,


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Cl2 is a pure element. To determine the charge on each Cl atoms, we divide the two bonded

electrons equally between the two Cl atoms, hence assigning 7 electrons to each Cl. Hence the

formal charge on each Cl is 0.

In a similar fashion we can determine the oxidation number of an atom bonded to another

atom. We can assign electrons to a bonded atoms, compare that number to the number in the

outer shell of the unbonded atoms, and see if there is an excess or lack. In these cases, the

same number of electrons get assigned to each atoms as when we are calculating formal

charge. Hence the oxidation numbers are equal to the formal charge in these examples. Clearly,

Na went from an oxidation number and formal charge of 0 to 1+ and Cl from 0 to 1-. Therefore,

Na was oxidized by the oxidizing agent Cl2, and Cl2 was reduced by the reducing agent Na.

Let's consider other similar redox reactions:

• 2Mg(s) + O2(g) 2MgO(s)

• Fe(s) + O2(g) Fe oxides(s)

• C(s) + O2(g) CO2(g)

In the first two reactions, a pure metal (with formal charges and oxidation numbers of 0) lose

electrons to form metal oxides, with positive metal ions. The oxygen goes form a formal charge

and oxidation number of 0 to 2- and hence is reduced.

What about the last case? Each atom in both reactants and products has a formal charge of 0.

This reaction, a combustion reaction with molecular oxygen, is also a redox reaction. Where are

the electrons that are lost or gained?


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This can be determined by assigning the electrons in the different molecules in a way slightly

different than we did with formal charge. For shared (bonded) electrons, we give both

electrons in the bond to the atom in the shared pair that has a higher electronegativity. Next

we calculate the apparent charge on the atom by comparing the number of assigned electrons

to the usual number of outer shell electrons in an atom (i.e. the group number). This apparent

charge is called the oxidation number. When we use this method for the reaction of C to CO2,

the C in carbon dioxide has an oxidation number of 4+ while the two oxygens have an oxidation

number of 2- . Clearly, the C has "lost electrons" and has become oxidized by interacting with

the oxidizing agent O2. as it went from C to CO2. If the atoms connected by a bond are identical,

we split the electrons and assign one to each atom. In water, the O has an oxidation number of

2- while each H atom has an oxidation number of 1+. Notice that the sum of the oxidation

numbers of the atoms in a species is equal to the net charge on that species.

What we have done is devise another way to count the electrons around an atom and the

resulting charges on the atoms. See the animation below to review electron counting, and the 3

"types of charges" - partial charges, formal charges, and now oxidation numbers.

Consider an O-X bond, where X is any element other than F or O. Since O is the second most

electronegative atom, the two electrons in the O-X bond will be assigned to O. In fact all the

electrons around O (8) will be assigned to O, giving it always an oxidation number of 2-. This will

be true for every molecule we will study this year except O2 and H2O2 (hydrogen peroxide). Now

consider a C-H bond. Since C is nearer to F, O, and N than is H, we could expect C (en 2.5) to be

more electronegative than H (en 2.1). Therefore, both electrons in the C-H bond are assigned to


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C, and H has an oxidation number of 1+. This will always be true for the molecules we study,

except H-H. A quick summary of oxidation numbers shows that for the molecules we will study:

• O always has an oxidation # of 2- (except when it is bonded to itself or F)

• H always has an oxidation # of 1+ (except when it is bonded to itself)

• The sum of the oxidation numbers on a compound must equal the charge on the

compound (just like the case of formal charges)

Notice in each of the reactions above, oxygen is an oxidizing agent. Also notice that in each of

these reactions, a pure element is chemically changed into a compound with other elements.

All pure, uncharged elements have formal charges and oxidation numbers of 0. When they

appear as compounds in the products, they must have a different oxidation number. The

disappearance or appearance of a pure element in a chemical reaction makes that reaction a

redox reaction.

Now lets consider a more complicated case - the reaction of methane and oxygen to produce

carbon dioxide and water:

CH4 + O2 H2O + CO2

Since H has an oxidation # of 1+, the oxidation # of C in CH4 is 4-, while in CO2 it is 4+. Clearly C

has been oxidized by the oxidizing agent O2. O2 has been reduced to form both products.

Now consider a series of step-wise reactions of CH4 ultimately leading to CO2

.


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You should be able to determine that the oxidation numbers for the central C in each molecule

are 4-, 2-, 0, 2+, and 4+ as you proceed from left to right, and hence represent step-wise

oxidations of the carbon. Stepwise oxidations of carbon by oxidizing agents different than O2

are the hallmark of biological oxidation reactions. Each step-wise step releases smaller amounts

of energy, which can be handled by the body more readily that if it occurred in "one step", as

indicated in the combustion of methane by O2 above.

You may have learned in a previous course that in oxidation reactions, there is an increase in

the number of X-O bonds, where X is some atom. Alternatively, it also involves the decrease of

X-H bonds. Reduction would be the opposite case - decreasing the number of X-O bonds and/or

increasing the number of X-H bonds. This rule applies well to the above step-wise example.

Redox reactions are common in nature. Some common redox reactions are those that occur in

batteries, when metals rust, when metals are plated from solutions, and of course combustion

of organic molecules such as hydrocarbons (like methane and gasoline) and carbohydrates (like

wood). A simple redox reaction that leads to the plating or deposition of a pure metal from a

solution of that metal is shown below.

Cu (s) Cu (aq) + 2e- (half equation/reaction)

Ag+ (aq) + e- Ag (s) (half equation/reaction)

Cu(s) + 2Ag+(aq) Cu2+(aq) + 2Ag(s)

In this reaction, pure silver metal - Ag(s) is plated on the surface of Cu(s) In this reaction:

• Cu is oxidized and is the reducing agent

• Ag+ is reduced and is the oxidizing agent


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If you look at the products, you could imagine they could also react in a reverse of the original

reaction to produce the original reactants.

Cu2+(aq) + 2Ag(s) Cu(s) + 2Ag+(aq)

In this reaction, pure copper metal - Cu(s) would be plated on the surface of the Ag(s). In this

reaction:

• Ag(s) is oxidized and is the reducing agent

• Cu2+ is reduced and is the oxidizing agent

Why doesn't this reverse reaction also occur? It actually does to a small extent. You could

actually envision the original reaction as reversible:

Cu(s)RA + 2Ag+(aq)OA Cu2+(aq)OA + 2Ag(s)RA

Where RA indicates which reactants/products are potential reducing agents in the forward and

reverse reactions, and OA indicates potential oxidizing agents. Which way does this reaction

go? Suffice it to say, the reaction goes in the direction from the strongest oxidizing and reducing

agents to the weakest. In the above example, Cu(s) is the stronger RA and Ag+ is the stronger

OA.

In redox, reactions, a pure element is either consumed or formed in the chemical

reaction. This would include all oxidation reactions involving O2. A reaction which

increases the number of bonds from an atom to oxygen represents an oxidation of that

atom and a reduction of oxygen. A decrease in the number of bonds from an atom to

oxygen represents a reduction of the atom. Both an oxidizing and reducing agent must

be present.


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Microbial Redox Process

Important redox reactions that are carried out by microorganisms are summarized here:

The notation [CH2O] is used to denote a fragment of an arbitrary carbohydrate.

Photosynthetic Production of Biomass

Photosynthetic microorganisms (algae and some bacteria) carryout photosynthesis reactions, in these reactions, energy-rich carbohydrate molecules are produced by combining carbon dioxide and water, using energy derived from sunlight.

These reactions can be written:

From a Redox Perspective:

CO2 + H2O [CH2O] + O2

Carbon is reduced from oxidation state +4 to 0, and oxygen is oxidized from -2 to 0

Aerobic Respiration:

In the presence of oxygen, microorganisms degrade biomass to form carbon dioxide and water. Chemical energy that is released can be used by the organisms.

[CH2O] + O2 CO2 + H2O

This process is the reverse of photosynthesis, carbon is oxidized and oxygen is reduced.

Nitrogen Fixation

• In the atmosphere, nitrogen is almost entirely in the form of N2 and is in oxidation state 0.

• The nitrogen in biological system is mostly in the form of an amine –NH 2, Which is very closely related to ammonia (NH3) and ammonium ion (NH4

+) here nitrogen is in oxidation state -3.

• Nitrogen in water and soil is in form of nitrate (NO3-) in which nitrogen is in oxidation state +5.

Microorganisms play an essential role in the movement of nitrogen among these oxidation states.

Compounds such as ammonia and nitrate contain a single nitrogen atom as fixed nitrogen species.

Certain groups of bacteria are capable of converting gaseous nitrogen to fixed nitrogen, in the form of the ammonium ion.

Energy from the oxidation of biomass to CO2 is used to reduce the nitrogen in N2 to ammonium.


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3[CH2O] + 2N2 + 3H2O + 4H+ 3CO2 + 4NH4+

Nitrification

In the nitrification, nitrogen in the ammonium ion is oxidized from -3 to +5, with oxygen as oxidizer.

NH4+ + 2O2 NO3

- + 2H+ + H2O

Plants absorb nitrogen more efficiently in the form of nitrate than an ammonium, so redox reaction can enhance the effectiveness of ammonia-based agricultural fertilizers.

Nitrate Reduction or Denitrification

When oxygen is not available as the oxidizer to degrade biomass, microorganisms can use nitrate as the oxidizer (electron acceptor).

Nitrate Reduction: is used in some wastewater treatment systems to convert fixed nitrogen to N2 gas, which can then be safely released to the atmosphere. This process is called denitrification, since nitrogen is removed from the aqueous system.

Nitrogen in municipal wastewater begins in a reduced state (-3), the overall process involved two steps:

Nitrification in an aerobic reactor, followed by denitrification in an anaerobic reactor, four nitrogen atoms, being reduced from +5 to 0, can fully oxidize five carbon atoms from 0 to +4

NH4+ + 2O2 NO3

- + 2H+ + H2O (aerobic reaction)

5[CH2O] + 4NO3- + 4H+ 5CO2 + 7H2O + 2N2 (anaerobic reaction)

Sulfate Reduction:

Some environments that contain biodegradable materials lake both oxygen and nitrate to serve as the oxidizing agent, in such cases, sulfate may serve that role.

The conversion of one sulfur atom from +6 in sulfate to -2, in hydrogen sulfide oxidizes two carbon atoms from 0 to +4 oxidation states.

2[CH2O] + 2H+ + SO42- 2CO2 + 2H2O + H2S

This reaction can occur in stagnant anaerobic marine sediments that are supplied with decaying biomass, algae or seaweed accumulation.

If the rate of biomass accumulation is high, oxygen can be rapidly depleted from the sediments.

Methane Formation (Methanogenesis):


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In the absence of oxygen, nitrate, and sulfate, biomass can still be converted to carbon dioxide as:

2[CH2O] CO2 + CH4

This is an interesting redox reaction, since the two carbon atoms begins in oxidation state zero (0). One carbon atom is oxidized to +4, and the other is reduced to -4.

Methane generation process is exploited in seawater treatment to convert excess microbiological material to gases, which are more easily handled for disposal.

Table-Oxidation States of Some Chemical Elements

Element Oxidation state

Species Formula

Sulfur

-2 Hydrogen sulfide H2S0 Elemental sulfur S

+4 Sulfur dioxide SO2

+6 Sulfate ion SO42-

Carbon

-4 Methane CH4

0 Soot, graphite C+2 Carbon monoxide CO+4 Carbon dioxide CO2

Nitrogen

-3 Ammonia NH3

0 Nitrogen gas N2

+2 Nitric oxide NO+3 Nitrite ion NO2

-

+4 Nitrogen dioxide NO2

+5 Nitrate ion NO3-

Oxygen-2 Almost all compounds --1 Hydrogen peroxide H2O20 Oxygen gas O2

Hydrogen0 Hydrogen gas H2

+1 Hydrogen ion H+

Chlorine

-1 Chlorine ion Cl-

0 Chlorine gas Cl2

+1 Hypochlorous acid HOCl+7 Perchloric acid HClO4

Hints when trying to predict products:


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• First look for obvious precipitation reactions - two aqueous salt solutions mixed

together with one of the salts containing an ion that quite often is insoluble (such as Fe,

Ba, Ag, Hg, Pb)

• Then looks for other possible precipitation reactions in which one of the above ions is

present in an aqueous salt but the other solution is an acid which might provide an

anion (like sulfate or chloride) that could precipitate with the above ion.

• Next identify easy acid/base reactions - in which you have an easily identifiable acid

(nitric, sulfuric, hydrochloric) and a hydroxide salt as the base

• Now look for easily identifiable redox reactions in which O2 is a reactant and probably an

oxidizing agent and the other reagent is a hydrocarbon, a carbohydrate, a metal, etc.

The hydrocarbon and carbohydrate will react to form carbon dioxide and water, and the

metal will react to form oxides.

• Finally look for other redox reactions, such as when you have a pure metal interacting

with a ion of another metal or an acid which could dissolve it to form a salt of the metal.

Precipitation ReactionAn aqueous solution is one that is occurring in water. What makes water significant is that it can allow for substances to dissolve and/or be dissociated into ions within it.

Reaction Types

Several schemes have been developed to categorize chemical reactions. The one we will use most often is charge transfer: These reactions involve making and breaking bonds and the transfer of charge between substances. If a proton is transferred from one substance to another, the reaction is called an acid/base reaction. An electron transfer occurs in an oxidation/reduction reactions. Finally soluble ions of salts can be "transferred" towards each other in solution and form an insoluble solid. This reaction is called a precipitation reaction.

Properties of Precipitates


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Precipitates are the products of a precipitation reaction, in which certain cations and anions combine to produce and insoluble solid. The determining factors of the formation of a precipitate can vary: some depend on temperature, such as solutions used for buffers, while others are dependent only on solution concentration. The solids produced in precipitate reactions are crystalline. The solid can then be suspended throughout the liquid or it can fall to the bottom of the solution. The liquid that remains above the precipitate is called the supernatant liquid.

Here is a diagram of the formation of a precipitates in the solution.

Precipitation and Double Replacement Reactions

Most precipitation reactions that occur are almost always either single replacement reactions or double replacement reaction. Double replacement reactions make up the majority of precipitation reactions, and we will be using a double replacement reaction for this particular explanation. The equation for a double-replacement reaction is as follows:

AB + CD → AD + CB

As illustrated above, a double replacement reaction occurs when two ionic reactants dissociate and bond with the respective anion or cation from the other reactant. This can be thought of as "switching partners," that is, the two pairs pictured above "lose" their partner and form a bond with a different partner:


36


A double replacement reaction is specifically classified as a precipitation reaction when the chemical equation in question occurs in aqueous solution and one of the products formed is insoluble. An example of a double replacement reaction is as follows:

CdSO4(aq) + K2S(aq) → CdS(s) + K2SO4(aq)

As you can see, both of the reactants are aqueous and the one of the products is solid. Because the reactants are ionic, and the fact that they are aqueous, i.e. in water, means that these reactants will dissociate and thus are soluble. However, there are certain solubility rules that dictate that some ionic molecules are not insoluble in water. These molecules will form a solid that precipitates throughout the solution.

Solubility Rules

Whether an aqueous double replacement reaction forms a precipitate is dictated by the solubility rules. They are as follows:

• Salts formed with group 1 cations and NH4+ cations are soluble. There are some

exceptions for certain Li+ salts.

• Acetates (C2H3O2-), nitrates (NO3

-), and perchlorates (ClO4-) are soluble.

• Bromides, chlorides, and iodides are soluble

• Sulfates (SO42-) are soluble. Exceptions to this rule are sulfates formed with Ca2+,

Sr2+, and Ba2+.

• Salts containing silver, lead, and mercury (I) are insoluble.

• Carbonates (CO22-), phosphates (PO4

3-), sulfides, oxides, and hydroxides (OH-) are insoluble. Sulfides formed with group 2 cations and hydroxides formed with calcium, strontium, and barium are exceptions to the rule.

Make sure to follow the lower-numbered guidelines in case of conflict between two rules.

Do you understand these concepts so far? Here are some examples to help you find out. Complete the double replacement reaction and indicate the states of matter that the products are in each chemical equation.

Example 1:

1. NaOH(aq) + MgCl2(aq) →

Because double replacement reactions "switch partners," we can predict the products of this reaction regardless of their states of matter. That means that the Na+ cation bonds with Cl- anion, and the Mg2+ cation bonds with the (OH)- to create:

2NaOH + MgCl2 → 2NaCl + Mg(OH)2

Now we must use consult the solubility rules to determine whether or not these products are soluble. If we consult the rules, we find that both group 1 cations (Na+) and chlorides are soluble, so we can conclude that NaCl will be soluble in water. However, hydroxides are insoluble, and thus Mg(OH)2 is insoluble in water; that is, Mg(OH)2 is a precipitate. So the resulting equation is:

2NaOH(aq) + MgCl2(aq) → 2NaCl(aq) + Mg(OH)2(s)


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Example 2:

2.CoCl2(aq) + Na2SO4(aq) →

Based on our knowledge of double replacement reactions, we can predict that the products of this reaction are CoSO4 and NaCl. Upon looking at the solubility rules, we can determine that CoSO4

is soluble because sulfates are soluble. Similarly, we find that NaCl is soluble, because chlorides are soluble. Thus the resulting equation looks as follows (after balancing, of course):

CoCl2(aq) + Na2SO4(aq) → CoSO4(aq) + NaCl(aq)

This particular example is important because both the reactants and the products are aqueous. This means that no precipitation is formed. As we will see in the next section, the presence of a precipitate dictates whether a double replacement reaction in aqueous reaction goes into completion.

Net Ionic Equations

The formation of a precipitate is a very important component of a double replacement reaction that occurs in aqueous solution: it allows us to write what is called a net ionic equation. To better understand the definition of a net ionic equation, let's look back on the equation for the double replacement reaction:

AB + CD → AD + CB

Because this particular reaction is a precipitation reaction, we can assign states of matter to each variable pair.

AB(aq) + CD(aq) → AD(aq) + CB(s)

The first step to writing a net ionic equation is to separate the soluble (aqueous) reactants and products into their respective cations and anions. Precipitates, as we know, do not dissociate in water, so we leave the solid in the reaction alone. The resulting equation would look something like:

A+(aq) + B-(aq) + C+(aq) + D-(aq) → A+(aq) + D-(aq) + CB(s)

If we look at this equation, we notice that the cation A+(aq) and the anion D-(aq) are present on both sides. These ions are called spectator ions, because they do not contribute to the formation of the precipitate. In the formation of the net ionic equation, we cancel out these spectator ions:

A+(aq) + B-(aq) + C+(aq) + D-(aq)→ A+(aq) + D-(aq) + CB(s)

The net ionic equation would then be:

C+(aq) + B-(aq) → CB(s)

It is important to note that the net ionic equation is reliant on the formation of a precipitate. If both products are aqueous, a net ionic equation cannot be written because all ions are considered to be spectator ions. Therefore, no reaction occurs.

Dissolution and Precipitation of Solids

Table-Solubility Products for Some Ionic Solids (at T=25oC)


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Compound Equilibrium relationship Ksp

Aluminum hydroxide Al(OH)3 ↔ Al3+ + 3OH- 1 X 10-32 M4

Cadmium hydroxide Cd(OH)2 ↔ Cd2+ + 2OH- 2 X 10-14 M3

Calcium carbonate CaCO3 ↔ Ca2+ + CO32- 5 X 10-9 M2

Calcium fluoride CaF2 ↔ Ca2+ + 2F- 3 X 10-11 M3

Calcium hydroxide Ca(OH)2 ↔ Ca2+ + 2OH- 8 X 10-6 M1

Calcium phosphate Ca3(PO4)2 ↔ 3Ca2+ + 2PO43- 1 X 10-27 M5

Calcium sulfate CaSO4 ↔ Ca2+ + SO4 2- 2 X 10-5 M2

Chromium(III) hydroxide Cr(OH)3 ↔ Cr3+ + 3OH- 6 X 10-31 M4

Iron(II) hydroxide Fe(OH)2 ↔ Fe2+ + 2OH- 5 X 10-15 M3

Iron(III) hydroxide Fe(OH)3 ↔ Fe3+ + 3OH- 6 X 10-38 M4

Magnesium carbonate MgCO3 ↔ Mg2+ + CO32- 4 X 10-5 M2

Magnesium hydroxide Mg(OH)2 ↔ Mg2+ + 2OH- 9 X 10-12 M3

Nickel hydroxide Ni(OH)2 ↔ Ni2+ + 2OH- 2 X 10-16 M3

Practice Problems:

Here are a few practice problems to help you determine whether you understand the material. From the given information, write the net ionic equation for the reaction, including states of matter. Highlight the text after the number under the solutions section for the solutions.

• Fe(NO3)3(aq) + NaOH(aq) →

• AlSO4(aq) + BaCl2(aq) →

• HI(aq) + Zn(NO3)2(aq) →

• CaCl2(aq) + Na3PO4(aq) →

• Pb(NO3)2(aq) + K2SO4 (aq) →

Solutions and Colloids

All chemical substances may be classified as elements, compounds, or mixtures.

Mixtures:

Mixtures are similar to compounds in that both can be broken down into simpler substances. But mixtures and compounds are different:

The ratio of weight of the simpler substances may vary in mixtures, but the ratio is always the same in compounds.

For example, the weight ratio of hydrogen to oxygen in water is always the same; however the ratio of alcohol to water varies in mixtures.

Solutions are mixtures because their composition may vary.

A solution is homogeneous mixture on an ionic or molecular scale of two or more different substances.


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Homogeneous: means made up of similar parts. In homogeneous mixtures the substances are uniformly mixed in one another.

Heterogeneous: made up of dissimilar parts. In heterogeneous mixtures the distribution of substances is not uniform.

For example, if fresh raw milk is allowed to stand for a period of time, the cream rises to the surface in a clearly detectable layer. Raw milk is heterogeneous mixture. Including tea, coffee, coke, seawater and sweets are the examples.

Colloids

The primary difference between colloids and solutions is the size of their particles.

In colloids the dispersed particles are large. Particles reflect light.

Colloids are quite common, such different substances as milk, shaving cream, smoke, and fog are all colloids.

Colloid such as milk is like a sugar solution in that the dispersed (dissolved) substances cannot be separated out by filtration. A suspension such as sand in water can be separated by filtration.

Factors Affecting Solution

The amount of the solute that will dissolve in a specific amount of a given solvent is called the solubility of the solute.

The primary factors affecting on the solubility of one substance in another are the nature of the solute, and the solvent, the temperature, and the pressure of the gas (in solution of gases in liquids).

The nature of the solute and solvent

A rule of thumb about solubility is “likes dissolve in likes”.

That is, polar solutes dissolve in polar solvents and non polar solutes dissolve in non polar solvents.

Example: NaCl in H2O = solution of sodium chloride in water (polar)

Oil dissolve in gasoline (non polar)

Temperature:


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For most common solution of solids in liquids, the solubility of the solute in the solvent generally increases as the temperature increases.

For example: the solubility of table sugar (sucrose) in water is:

179 grams per 100 ml of H2O at 0 oC and

487 grams per 100 ml of H2O at 100 oC

Exception: the solubility of gases in liquids decrease with increased temperature.

Pressure:

Pressure has a significant effect on the solubility of gases in liquids.

Henry’s Law: states that the solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid.

Thus, if the pressure of gas above a liquid is doubled, its solubility doubled.

The factors affecting the Rate of Dissociation

In general, the rate with which a given solute will dissolve in a given solvent is determined by the size of the solute particles, the rate of stirring, and the temperature.

Size of the solute particles:

The more finely divided solute, the more rapidly it will dissolve, the larger the surface area of the solute exposed to the solvent, the more rapidly it dissolves. The surface area of a substance increases very rapidly as it is subdivided into smaller particles.

Stirring:

The more vigorous the stirring, the more quickly the solvent in contact with the solute is changed, thus the more rapid the rate of dissociation.

Temperature:

At higher temperature both the solute and solvent particles have, greater energy of motion. Because of their increased kinetic energy, the solute particles break away from one another more rapidly.


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• Solution Concentration

Qualitative and Percent Methods:

• In using a solution you often need to know the amount (weight or volume) of solute in the solution

The reaction between many substances (acids and bases, for example) takes place in solution.

• When you are dealing with the chemical properties of the solution, you need to know the amount of solute in a given amount of the solution.

Also, when solutes are dissolved in liquids, the presence of the solute affects the physical properties of the liquid.

For example, when sugar is dissolved in water, the boiling point of the solution is higher and the freezing point of the solution is lower than for pure water.

The change in the properties of the liquids depends on the ratio of solute to solvent molecules.

• So when you are considering some of the physical properties of solutions, you need to know the amount of solute per amount of solvent.

• Dilute Solution: contains relatively small amount of solute.

• Concentrated Solution: contains relatively large amount of solute.

• Saturated Solution:

-That contains the amount of dissolved solute necessary for there to be equilibrium between dissolved and un-dissolved solute.

-Its concentration depends on the solubility of the solute, the temperature, and the pressure for the gas.

-A saturated solution is not necessarily a concentrated solution. Saturated solutions of relatively insoluble substances contain a very small amount of solute. They are therefore, relatively dilute. On other hand, saturated solutions of soluble substances are quite concentrated.

• Supersaturated Solution:

For most substances a hot saturated solution contains more solute than a cold saturated solution of the substance. The solution will contain more solute than a saturated solution at that cooler temperature.

• Percent concentration:


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-There are three ways in which percent is used to express concentration.

To avoid confusion, specify ‘percent by weight”, “percent by volume”, and so on.

-One of the three general methods used to express the concentration of solutions in terms of the quantity of solute per amount of solution.

-Percent concentration expresses the number of parts of solute per 100 parts of solution.

• Percent by Weight

-Solution concentration in a percent is to indicate the weight of solute per 100 grams of solution.

Thus, a 5 percent solution contains 5 grams of solute per 100 grams of solution.

-Each milliliter (cubic centimeter) of water weights 1 gram.

Percent by weight = weight of solute/weight of solution X 100

• Percent by Volume: The concentration solution, of one liquid in another is often expressed as percent by volume.

Percent by Volume = volume of solute / volume of solution X 100

• Percent by Weight-Volume: In medicine, concentration is often expressed in terms of mass of solute per volume of solution.

mg % = mg of solute / ml of solution X 100

For example,

The concentration of glucose (sugar) in the blood, expressed in terms of milligrams of glucose per 100 ml of blood.

-There are many substances dissolved in the blood, and the concentrations of these substances normally vary a lot from time to time. So the weight of a given volume of blood fluctuates almost continuously. For this reason the weight of a given substance such as glucose per reference VOLUME of blood rather than per reference WEIGHT of blood is used.

Percent by Weight Calculation:

Problem: Calculate the percent by weight concentration for a solution that contains 9.00 g of sucrose (table sugar) in 30.0 g of solution.

Solution: Percent sucrose = 9.00 g sucrose / 30.0 g solution X 100

= 0.300 g sucrose / 1.00 g solution X 100 = 30.0 percent


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Problem: How would you prepare 500 g of 12.0 percent NaCl solution?

(Note: A 12.0 percent solution contains 12.0 g of NaCl per 100 g of solution)

Solution: g NaCl = 12.0 g NaCl / 100 g solution

= 0.120 g NaCl / 1.00 g solution X 500 g solution = 60.0 NaCl

Wt. H2O = 500 g – 60.0 = 440 g.

Therefore, to prepare 500 g of 12.0 percent NaCl solution, you would weight out 60.0 g NaCl and dissolve it in 440 g (440 ml) of water.

Problem: How many grams of sugar must be added to 500 g (500 ml) of water to prepare a 15.0 percent solution? (100 g of solution will contain 15.0 g of sugar and 85.0 g of water).

Solution: g sugar = 15.0 g sugar / 85.0 g H2O X 500 g H2O = 88.2 g sugar.

Problem: How many grams of water must be added to 16.0 g of sugar to give a 4.00 percent solution? (100 g of solution will contain 4.00 g of sugar and 96.0 g of H2O)

Solution: g H2O = 96.0 g H2O / 4.00 g sugar X 16.0 g sugar = 384 g H2O

• Solution Concentration

Moles of solute per Volumes of solution:

-The concentration of solute in a solution in terms of percent tells simply the solution is dilute or concentrated. Molarity tells how many molecules of solute are involved.

-However, chemical reaction occurs between molecules of the solutes.

-The weight of a mole (Avogadro’s number of molecules) of one substance may be quite different from the weight of a mole of another substance.

For example,

-A mole of sodium hydroxide (NaOH – molecular wt. = 40.0) weights 40.0 grams and a mole hydrobromic acid (HBr – molecular wt. = 80.9) weights 80.9 grams.

-A 10 percent solution of NaOH and a 10 percent of HBr contain the same weight of solute per 100 grams of solution. But the number of molecules of solute in each solution is very different.

-Expressing concentration in terms of the moles of solute per volume of solution is usually more useful in chemistry than concentration expressed as percent.


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Molarity (M) M = no. of moles of solute / 1 liter of solution

-The molarity expresses the number of moles of solute per unit volume of solution-not volume of solvent or weight of solution.

-The volume of solvent depends on both the molarity of the solution and the density (weight per unit volume) of the solute.

The following relationships are used in solving problems dealing with molar concentration:

• 1 mole = 1 gram-molecular weight (g-molecular wt.)

• Molarity = M = no. of moles of solute / no. of liters of solution

= no. of moles of solute / 1 liter of solution

Problem: What is the molarity of a solution if 5 liter of it contains 4 moles of solute?

Solution: M = 4 moles of solute / 5 liter of solution

= 0.8 mole of solute / 1 liter of solution = 0.8 M

• No. of moles of solute = M X no. of liters solution

= moles of solute / 1 liter of solution X no. of liters of solution

M (moles of solute in 1 liter) times the number of liters of a solution gives the number of moles of solute in that volume of the solution.

Problem: How many moles of solute are in 0.30 liter of 1.5 M solution?

Solution: moles solute = 1.5 moles / 1 liter X 0.30 liter = 0.45 mole solute

• No. of moles of solute = no. of grams of solute

X 1 mole of solute / 1 g-molecular wt. of solute

Problem: How many moles of NaOH (molecular wt. = 40.0) are in 100 g of the compound?

Solution: moles NaOH = 100 g NaOH X 1 mole NaOH / 40.0 g NaOH = 2.5 moles NaOH

• No. of grams of solute = no. of moles of solute


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X g-molecular wt. of solute / 1 mole of solute

Problem: What is the weight in grams of 0.25 moles of NaOH (molecular wt. = 40.0)?

Solution: g NaOH = 0.25 mole of NaOH X 40.0 g NaOH / mole NaOH = 10 g NaOH.

Chemical Equivalency

Molar concentration has its limitations because all molecules do not react on a one-to-one basis.

For example,

The equations for the reaction of NaOH solution with hydrochloric acid, sulfuric acid, and phosphoric acid:

NaOH + HCl NaCl + H2O

2NaOH + H2SO4 Na2SO4 + 2H2O

3NaOH + H3PO4 Na3PO4 + 3H2O

1 mole of NaOH equal to 1 mole of HCl

2 moles of NaOH equals to 1 mole of H2SO4

3 moles of NaOH equals to 1 mole of H3PO4

Normality (N)

The number of equivalents (gram-equivalent weight) of solute per liter of solution

Thus, one normal (1N) solution contains 1 gram-equivalent weight of solute per liter of solution

Normality (N) = no. of equivalent of solute / 1 liter of solution

Equivalent Weight

The equivalent weight of a substance is the molecular weight of the substance divided by its equivalency per mole in the particular reaction.


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Substance Molecular weight Equivalency Per mole

Equivalent weight

NaOH 40.0 1 40.0 / 1 = 40.0HCl 36.5 1 36.5 / 1 = 36.5H2SO4 98.0 2 98.0 / 2 = 49.0H3PO4 98.0 3 98.0 / 3 = 32.7

Calculations related to Normality

The following relationships are important when making calculations related to normal concentration:

• N = no. of equivalents of solute / no. of liters of solution

= no. of eq. of solute / 1 liter of solution

• 1 eq. = 1 g-eq. wt.

• 1 g-eq. wt. = g-molecular wt. / equivalency per mole of solute

= g-molecular wt. X 1 mole / no. of eq. per mole of solute

(Remember, equivalency per mole equals the number of protons an acid molecule yields or a base molecule accepts in the reaction under consideration.)

• No. of eq. (g-eq. wt.) = no. of g of solute X 1 eq. of solute / g-eq. wt. of solute

• No. of eq. (no. of g-eq. wt.) of solute = N X no. of liters of solution

= eq. of solute / liter X no. of liters.

• No. of g of solute = no. of eq. of solute X g-eq. wt. of solute / eq. of solute

• N = M X no. of eq. of solute / 1 mole of solute

M = N X 1 mole of solute / no. of eq. of solute

Problem: Calculate the normality of a solution that contains 2.8 equivalents of solute in 1.4 liter of solution.

Solution: N = 2.8 eq. solute / 1.4 liter solution = 2.0 eq. solute / 1 liter solution = 2.0 N

Problem: Calculate the number of equivalents of solute in 500 ml of 6.0 N solution

Solution: eq. solute = 6.0 eq. / liter solution X {500 ml X 1 liter / 1000 ml} = 3.0 eq.


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Problem: A solution was prepared by dissolving 101 g of Ca (OH)2 (molecular wt. = 74.1) in enough water to yield 5.00 liters of solution. What is the normality of the solution? [Assume that each Ca(OH)2 molecules accept two protons in reaction].

Solution:

First, calculate the gram-equivalent weight of Ca (OH)2.

g-eq. wt. Ca(OH)2 = 74.1 g Ca(OH)2 / mole Ca(OH)2 X 1 mole Ca(OH)2 / 2 eq. Ca(OH)2

= 37.0 g Ca(OH)2 per eq.

Than

N = 101 g ca(OH)2 / 5.0 liter solution X 1 eq. Ca(OH)2 / 37.0 g Ca(OH)2

= 0.55 eq. Ca(OH)2 / 1 liter solution = 0.55 N

Problem: A solution was prepared by dissolving 6.0 g of NaOH (molecular wt. = 40) in 200 ml of solution. Calculate the normality of the solution.

Solution: g-eq. wt. NaOH = 40 g NaOH / mole NaOH X 1 mole NaOH / 1 eq. NaOH = 40 g NaOH / eq.

(NaOH has only one OH- per molecule, so it can accept one proton per molecule-or 1 mole of NaOH can accept 1 mole of protons).

N= 6.0 g NaOH / 200 ml solution X 1 eq. wt. NaOH / 40 g NaOH X 1000 ml solution / 1 liter of solution

= 0.75 eq. wt. / 1 liter = 0.75 N

Problem:

How would you prepare 1.2 liter of 2.4 N H2SO4 solution? (Assume that H2SO4 donates both protons when it reacts)

Solution:

g-eq. wt. H2SO4 = 98 g H2SO4 / mole H2SO4 X 1 mole H2SO4 / 2 g-eq. wt

= 49 g H2SO4 / eq. wt. H2SO4

g H2SO4 = [2.4 eq. H2SO4 / 1 liter solution X 1.2 liter solution] X 49 g H2SO4 / eq. wt. H2SO4

Therefore, you would dissolve 140 g of H2SO4 in enough H2O to make 1.2 liter solution.


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Problem:

What volume of 0.900 N H3PO4 solution can be prepared from 147 g of H3PO4 (molecular wt. = 98.0)? (Assume that all three hydrogen ions of H3PO4 are replaced).

Solution:

g-eq. wt. H3PO4 = 98.0 g H3PO4 / mole H3PO4 X 1mole H3PO4 / 3 g-eq. wt.

= 32.7 g H3PO4 / g-eq. wt.

N = no. of g-eq. wt. / liter of solution

Therefore:

Liter solution = no. of g-eq. wt. / N

= no. of g-eq. wt. / eq. per 1 liter = no. of g-eq. wt. X 1 liter / eq.

1 H3PO4 solution = [147 g H3PO4 X 1 g-eq. wt. H3PO4 / 32.7 g H3PO4] X 1 liter / 0.900 eq.

= 5.00 liter solution

Solvents and SorbentsDevelopment of more efficient solvents is crucial in reducing the cost of carbon dioxide absorption from combustion exhaust gases. Several solvents have been proposed for carbon dioxide absorption, but the question is how the best candidate could be selected. Alkanolamines such as monothanolamine (MEA), diethanolamine (DEA), di-isopropanolamine (DIPA) and methyldiethanolamine (MDEA) have been the most common solvents used over the years for absorption process in removal of acid gases, CO2 and H2S from industrial and combustion exhaust gas streams. However, these alkanolamines are still deficient for carbon dioxide absorption due to inherent problems associated with their use in CO2 capture process. Different factors affect the efficacy of a solvent for carbon dioxide absorption, these include solvent solubility, vapor pressure, molecular weight and foaming tendency, degradation and corrosion properties; others are reaction kinetics, heat of reaction and regeneration energy requirement as well as the cyclic capacity. Environmental and cost factors are also to be considered.

Chemical Solvents

The most commonly used technology today for low concentration CO2capture is absorption with chemical solvents. This chemical absorption process is adapted from the gas processing industry where amine-based processes have been used commercially for the removal of acid gas impurities from process gas streams. However, problems of scale, efficiency, and stability become barriers when chemical solvents are used for high volume gas flows with a relatively smaller fraction of valuable product. The processes require large amounts of material undergoing significant changes in conditions, leading to high investment costs and energy consumption. In


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addition, degradation and oxidation of the solvents over time produces products that are corrosive and may require hazardous material handling procedures.

Amine Solvents

The currently preferred chemical solvent technology for carbon capture is amine-based chemical absorbent. CO2 in the gas phase dissolves into a solution of water and amine compounds. The amines react with CO2 in solution to form protonated amine (AH+), bicarbonate (HCO3-), and carbamate (A CO2 -) [9]. As these reactions occur, more CO2 is driven from the gas phase into the solution due to the lower chemical potential of the liquid phase compounds at this temperature. When the solution has reached the intended CO2 loading, it is removed from contact with the gas stream and heated to reverse the chemical reaction and release high-purity CO 2. The CO2-lean amine solvent is then recycled to contact additional gas. The flue gas must first be cooled and treated to remove reactive impurities such as sulfur, nitrogen oxides, and particulate matter.

Otherwise, these impurities may react preferentially with the amines, reducing the capacity for CO2, or irreversibly poisoning the solvent. The resulting pure CO2 stream is recovered at pressures near atmospheric pressure. Compression, and the associated energy costs, would be required for geologic storage.

Alkanolamines, simple combinations of alcohols and ammonia, are the most commonly used category of amine chemical solvents for CO2 capture. Reaction rates with specific acid gases differ among the various amines. In addition, amines vary in their equilibrium absorption characteristics and have different sensitivities with respect to solvent stability and corrosion. Alkanolamines can be divided into three groups :

• Primary amines, including monoethanol amine (MEA) and diglycolamine (DGA)

• Secondary amines, including diethanol amine (DEA) and diisopropyl amine

(DIPA)

• Tertiary amines, including triethanol amine (TEA) and methyldiethanol amine

(MDEA)

MEA, relatively inexpensive and the lowest molecular weight, is the amine that has been used extensively for the purpose of removing CO2 from natural gas streams. MEA has a high enthalpy of solution with CO2, which tends to drive the dissolution process at high rates. However, this also means that a significant amount of energy must be used for regeneration. In addition, a high vapor pressure and irreversible reactions with minor impurities such as CO2 and CS2 result in solvent loss.

Research on improved chemical solvents seeks a high absorption capacity for CO2 without a corresponding large energy requirement for regeneration. Other desirable properties include high chemical stability, low vapor pressure, and low corrosiveness. It has been shown that solvents based on piperazine-promoted K2CO3 can have reaction rates approaching that of MEA, but currently have lower capacity. Sterically hindered amines have been developed with similar capacity and possibly less regeneration energy requirement than conventional MEA absorbents. These modified amines attempt to balance good absorption and regeneration characteristics


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under some conditions due to the reduced chemical stability of the amine- CO2 anion [9]. Controlled species selectivity is also possible with these compounds.

Physical Absorption

Absorbents allow a gas to permeate a solid or liquid under one set of conditions, and desorb under others. The rate of absorption or desorption is temperature and pressure dependent. Smaller differences in conditions require less energy, but require more absorbent to capture CO2

at an equivalent rate.

Physical Solvents

Absorption in most current physical solvent systems occurs at high partial pressure of CO2 and low temperatures. The solvents are then regenerated by either heating, pressure reduction, or a combination of both. The interaction between CO2 and the absorbent is weak relative to chemical solvents, decreasing the energy requirement for regeneration. Capacity can be higher than chemical solvents, since it is not limited by the stoichiometry of the chemical system.

Physical solvent scrubbing of CO2 is well established. Selexol, a liquid glycol-based solvent, has been used for decades to process natural gas, both for bulk CO2 removal and H2S removal. Glycol is effective for capturing both CO2 and H2S at higher concentration. However, the CO2 is released at near atmospheric pressure, requiring recompression for transportation and geologic storage.

The Rectisol process, based on low-temperature methanol, is another physical solvent process that has been used for removing CO2. Glycerol carbonate is interesting because of its high selectivity for CO2, but it has a relatively low capacity.

Mixed Chemical-Physical Solvents

Some CO2 capture applications benefit from a mixture of physical and chemical solvents.

The most commonly used examples are Sulfinol, a mixture of the physical solvent sulfolane and the amines DIPA or MDEA, and Amisol, a mixture of methanol and secondary amines. These hybrid solvents attempt to exploit the positive qualities of each constituent under special conditions.

Physical Adsorption

Physical adsorption relies on the affinity of CO2 to the surface of a material under certain conditions without forming a chemical bond. Adsorbents can separate CO2 from a stream by preferentially attracting it to the material surface at high pressures through weak interactions such as van der Waals forces. During capture, the chemical potential of the adsorbed CO 2 is lower than the chemical potential of CO2 in the gas mixture.

Regenerable Physical Adsorbents

Regenerable adsorbents must have the ability to reverse the chemical potential of the adsorbed phase upon changing the conditions to remove the CO2. This is done primarily through changes in pressure or stripping with an easily separable gas such as steam.

Limited temperature changes can improve efficiency, but take time cycle due to the heat capacity of the adsorbent material. Since adsorption is a surface phenomenon, a successful adsorbent will have a high surface area to volume ratio. The central advantage of physical adsorption methods


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is the possibility for low energy requirement to regenerate the sorbent material and the quick regeneration time associated with changing the pressure.

Proposed absorbents include activated carbon, zeolites (molecular sieves), and promoted hydrotalcites. Current zeolite systems can produce nearly pure streams of CO2, but have high energy penalties due to vacuum pumps and dehumidification equipment. Hydrotalcites are most effective at high temperatures (450-600 K), enabling capture inside or near combustion or gasification chambers . Research is required to decrease the pressure difference requirement and increase the capacity of current adsorbents.

Membrane Separation Processes

Membrane systems include thin barriers that allow selective permeation of certain gases, allowing one component in a gas stream to pass through faster than the others. Membrane separation can be considered a steady-state combination of adsorption and absorption. A successful membrane allows the desired gas molecule to adsorb to the surface on one side, often at higher pressure. The molecule then absorbs into the membrane interior, eventually reaching the other side of the membrane where it can desorb under different conditions, such as low pressure.

Membrane gas separation processes have been widely used for hydrogen recovery in ammonia synthesis, removal of CO2 from natural gas, and nitrogen separation from air. Each of the membranes used in these capacities could be applied to carbon capture. Commonly used membrane types for CO2 and H2 separation include polymeric membranes, inorganic microporous membranes, and palladium membranes.

Polymeric membranes, including cellulose acetate, polysulfone, and polyimide are the most commonly used for separation of CO2 from nitrogen, but have relatively low selectivity to other separation methods. Inorganic membranes, able to withstand high temperatures, are capable of operating inside combustion or gasification chambers. Membrane reactors based on inorganic membranes with palladium catalyst can reform hydrocarbon fuels to mixture of H2 and CO2 and at the same time separating the high-value H2. Combining membranes with chemical solvents has also been proposed. Despite an extra energy requirement, this arrangement may eliminate problems associated with direct contact between the liquid solvent and gas mixture.

Most membranes have inherent difficulty achieving high degrees of gas separation due to varying rates of gas transport. Stream recycling or multiple stages of membranes may be necessary to achieve CO2 streams amenable to geologic storage, increasing energy consumption. However, the potential for high surface area could reduce the chemical potential difference required to drive gas separation.

Chemisorption

Gas molecules can chemically bond to the surface of some materials. The process is called associative if the molecule bonds in whole to the surface and dissociative if the gas molecule breaks up in order to form a bond. Chemisorbents are often composed of an active surface layer supported by an inert substrate. Proposed systems use small particles as substrates in order to provide large surface area. Regeneration drives the chemical reaction in reverse, often at elevated temperature.

Metal Oxide Air Separation


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Air separation allows a pure stream of oxygen to react with the fuel, creating an effluent of only CO2 and water or other useful products such as hydrogen. It is often easier with current technology to separate CO2 from water or hydrogen than from nitrogen. Reactive metal exposed to air will oxidize rapidly. The metal oxidation reaction is highly exothermic (ΔHOX ~ -950 kJ/mol). The oxides can then be endothermically (ΔHRED ~ -150 kJ/mol) regenerated by exposing them to a high temperature reducing environment. The oxygen combines with the fossil fuel to form carbon oxides and varying amounts of hydrogen-containing species depending on the type of fuel. The chemistry and geometry of this separation has allowed recent small-scale studies to obtain a nearly 100% pure stream of oxygen to react with the fuel. Phase separation of water from the resulting effluent could produce a pure stream of CO2. This complete process is commonly called chemical looping separation.

Research in metal oxide air separation is focused on cost and the physical and chemical stability of the oxygen carriers over many cycles. The particles usually consist of a reactive oxide and a supporting inert oxide. While various oxygen carrier particles are under consideration, copper, iron, manganese, and nickel are the most promising reactive metals.

No large-scale demonstration has been performed, but models predict that a power system utilizing metal oxide air separation has significant advantages. The lower irreversibility associated with the regeneration step relative to conventional combustion add to the already low energy requirement of the inherent separation of CO2 from nitrogen. Energy analyses show the resulting overall energy penalty could be as low as 400 kJ/kg CO2 for a natural gas combined cycle plant, assuming idealized chemical stability of the oxygen carrier.

Dry Chemical Absorbents

Under some conditions, CO2 can undergo a reversible chemical reaction with a dry absorbent material. The chemical reaction can be reversed by changing the conditions, resulting in the release of pure CO2. Sodium carbonate supported on an inert particle has been proposed as such an absorbent. An exothermic reaction of sodium carbonate with CO2 and water held at 60 to 70ºC forms primarily sodium bicarbonate. The products must be heated to 120 to 200ºC to reverse the reaction. Lithium zirconate is also being investigated for its high capacity chemisorbtion separation of CO2 at high temperatures.

Chemical Bonding

Some gas separation technologies use materials that create stable, thermodynamically favored chemical bonds with a gas in a mixture or a gas in solution. These materials can either be endothermically regenerated and used in a loop much like the sorbent technologies or form stable waste materials to be stored.

CO2 Mineralization

Some minerals will undergo thermodynamically favorable reactions with CO2, separating it from a gas stream and forming a stable, chemically bonded product. Although most proposed cycles have problems with kinetics due to the relatively low enthalpy of reaction, separation and conversion to a stable storage medium is accomplished in one step. Open-ended cycles such as this have the advantage of not requiring regeneration. The reactant minerals can be considered a separate resource that provides the energy requirement for separation.


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One proposed method involves a reaction of CO2 with CaCO3, or limestone, and water to form calcium and bicarbonate ions. These ions can be deposited into the ocean, short circuiting the residence of carbon in the atmosphere. Another method proposes enhancing the otherwise slow mechanism of silicate weathering. A gas stream containing CO2 could react with magnesium silicate to form magnesium carbonate and pure silicate. The large volumes of material involved present significant challenges for transportation and handling.

Phase Separation

Below certain temperatures, gas molecules are moving slow enough to succumb to weak intermolecular forces. Depending on the partial pressure of other gases in a mixture, condensing gases will form a distinct phase with a composition different from that of the vapor that is easily separated.

CO2 Clathrate

Clathrates are a phase of water in which the hydrogen-bonded structures encapsulate “guest” molecules of gas. The preferential formation of CO2 clathrate over other fossil fuel conversion effluent gases could be used as a method to capture CO2. The CO2 clathrates could later be dissociated, producing a pure stream of CO2. Formation of CO2 hydrates occurs at about 140 atm and at temperatures near the freezing point of water.


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Environmental chemistry (notes)12IEEM BY Muhammad Fahad Ansari 12IEEM14

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