Objectives:

The world around us is composed almost entirely of compounds and mixtures of compounds: rocks, coal, trees, and human bodies are all complex mixtures of chemical compound in which different kinds of atom are bound together.

In isolated atom each electron is under the influence of the nuclear charge and the charges of all of the other electrons present. When atoms coming together forming compounds, a rearrangement of electrons occur such that electrons of one atom are influenced by the nucleus of another atom, thus electrons are shared or transferred. If this rearrangement produces an energetically stable condition then bond formation can occur.

Figure 6.1 Schematic of the bond formation between two hydrogen atoms.

Through the covalent bond, each hydrogen atom gets to ‘use’ an electron from other atom as well as its own, thus obtaining a filled outer level. The shared pair of electron arrangement in the bond can be represented in several different ways:

• using Lewis symbols: H : H ( here dots indicate electrons);
• by the chemical formula: H₂
• by the structural formula: H - H

Chemical formula gives the number and kind of atoms in a molecule.

Structural formula shows individual bonds in a molecule.

Table 6.1 Diatomic molecules.

Oxygen, O₂, has a double bond because oxygen atom has six outer electrons (2s²2p⁴) and it need two more electrons to complete the outermost level to eight electrons (octet rule).

Octet rule states that atoms tend to combine in such a way that each obtains a filled outermost level of EIGHT electrons upon forming a chemical bond.

Nitrogen, N₂, has five outermost electrons (2s²2p³) per atom. Thus, it requires three more electrons to fill the outermost level. Therefore, a nitrogen molecule is very strong. Because our bodies cannot break down the strong triple bond, we cannot use the nitrogen gas in the atmosphere although four-fifths of the atmosphere consists of N_2.

Covalent bond generally forms between of nonmetals. Here electrons are shared equally between identical atoms. This type of bond is also known as a nonpolar covalent bond. The type of unequal sharing of electrons is called a polar covalent bond.

Example: The polar covalent bond in a hydrogen chloride H - Cl creates a dipole. Hydrogen chloride is called a polar molecule because it contains a dipole.

Dipole is a pair of equal but opposite charges that are separated by a small distance.

Figure 6.2 Polarity of water molecule.

Molecules are called nonpolar if they contain no dipole.

Example: CO₂ is nonpolar molecule because two dipoles of this molecule point in opposite directions and cancel each other

Electronegativity is a measure of the ability of an atom to attract to itself the electrons in a chemical bond.

• In the periodic table, across the period, electronegativity increases because the nucleus becomes more and more positively charged. Down a group, it decreases because the electron in the outermost orbitals are farther away from the nucleus of the atom. Thus the most electronegative elements are the nonmetals at the upper right of the periodic table. Fluorine is the most electronegative element, followed by oxygen and chlorine.

Bound energy is the amount of energy required to pull the bonded atoms apart.

Bond length is the distance between the atoms in a molecule.

• When metals form bonds with nonmetals, the differences in electronegativity between the atoms is so great that it become difficult to describe the bonding as simple as sharing of electrons. While in covalent bonds electrons of individual atoms in a molecule are considered to be shared between atoms, ionic bonds form when one or more electrons from one atom are completely transferred to a second atom.

Ionic bonds results from the attraction between appositely charged ions, namely cation and anions.

Ions are electrically charged atoms. Cation is a positively charged ion. A cation consists of a positive nucleus surrounded by electrons that are too few to balance the changes of all its protons. Cations are formed by a process called oxidation, which means the loss of one or more electrons.

Anion is a negatively charged ion. An anion consists of a positively nucleus surrounded by more than enough electrons to balance the charge of its protons. Anions are formed by reduction, which means the gain of one or more electrons.

• The elements most easily oxidized are the metals. The more metallic an element (the lower its electronegativity), the easier its atoms oxidized to become cations.
• The elements reduced more easily are the nonmetals. The more nonmetallic an element (the greater its electronegativity), the easier it is reduced to form an anion.

Table 6.2 Examples of metal elements.

Table 6.3 Examples of nonmetal elements.

Example of ionic bond formation:

Sodium chloride, NaCl, consists of the cation Na⁺ and the anion Cl^-

Sodium is oxidized 1s²2s²2p⁶3s¹ 1s²2s²2p⁶

losing an electron. sodium atom, Na sodium ion, Na⁺

Chlorine is reduced 1s²2s²2p⁶3s²3p⁵ 1s²2s²2p⁶3s²3p⁶

gaining an electron chlorine atom, Cl chloride, Cl^-

The sodium cation and the chloride anion have opposite charges and therefore attract each other.

Magnesium chloride, MgCl_2, consists of the cation Mg²⁺ and two anions Cl^-

Magnesium is oxidized losing two electrons:

1s²2s²2p⁶3s² 1s²2s²2p⁶

magnesium atom, Mg magnesium ion, Mg²⁺

Chlorine is reduced gaining one electron. Therefore, two chlorine atoms are needed for each magnesium atom.

Polyatomic ions are groups of atoms that are covalently bonded and have a positive or negative charge.

Example: The hydroxide ion, OH^-, is an ion composed of one hydrogen atom and one oxygen atom. OH^- goes into bonding as an ion with charge 1^- (for instance, NaOH sodium hydroxide).

NH₄⁺ ammonium ion; NO₃^- nitrate ion; NO₂^- nitrite ion; CO₃^2- carbonate ion

SO₃^2- sulfite ion; SO₄^2- sulfate ion; HSO₄^- hydrogen sulfate ion (bisulfate ion)

The term valence describes the number of electrons which an atom utilizes in bonding.

Oxidation number (or state) is defined as the charge on the atom in a given compound or ion, after bonding electrons have been allocated according to the general rules.

1. The oxidation number of hydrogen is +1 (except in metallic hydrides when it is -1 and in hydrogen gas when it is 0).

For instance, it is +1 in hydrochloric acid HCl, ammonia NH_3, water

H₂O, methane CH_4, but it is -1 in LiH

2. The oxidation number of fluorine is always -1 and it never forms more than one bond. Similarly, the oxidation number of all halogens is -1 (except when they are combined with more electronegative elements).

For instance, it is -1 in HI, but it is +1 in ICL.

3. The oxidation number of oxygen is -2 (except in peroxides, for instance H₂O_2, when it is -1 and in oxygen gas O₂ when it is 0).

For instance, it is -2 in CO, CO₂, SO₂, SO₃.

4. The oxidation number of an atom in an element is 0.

For instance, it is 0 in Na(s), Hg(l), O₂(g), O₃(g), N₂(g).

5. The oxidation number of monatomic ions is the charge of the ion

For instance, Na⁺ --> +1; Cl^- --> -1

6. In any molecule or ion, the sum of the oxidation numbers of all atoms present is equal to the net charge on molecule or ion. If there is no any charge the sum of oxidation numbers is equal to 0.

For instance, the sum of oxidation numbers for the hydrogen and oxygen in water, H₂O, is 0.

For instance, the sum of the oxidation number for the carbon and oxygen atoms in CO₃^2- is -2 because the oxidation state for O is -2 and for C is +4.

 

• The oxidation numbers (states) are not actual charges. Actual charge is given as n^- or n⁺, in turn oxidation number is given as -n or +n. For example, for NO₃^-, the overall charge is 1^-; the oxidation state for N is +5 and for O is -2.
• The concept of oxidation states (numbers) provides a way to keep track of electrons in oxidation-reduction reactions which are very important in atmospheric chemistry (see Lecture 7).

• Some elements (such as nitrogen, carbon, sulfur) can exist in the environment in a wide range of oxidation states.

Table 6.4 Oxidation states of nitrogen and some typical compounds.

*Compounds with -2 and -1 states are not normally found in the natural environment.

Chemical reaction is the transformation of two or more interacting reactants through the intermediary phase of an activated complex, in which existing chemical bonds are broken and new bonds are formed, in which the chemical bonds define the species before and after the reaction.

Chemical equation is representation of a chemical reaction showing the relative numbers of reactants and products.

For instance, A + B => C + D

where A and B are the reactants; C and D are the products.

NO + O₃=> NO₂ + O₂ (reactants: NO and O₃; products: NO₂ and O₂)

NO₂ + hn => NO + O (reactants: NO₂; products: NO and O)

• In a chemical reaction atoms are neither created nor destroyed.
• Atoms, and therefore mass, are conserved in a chemical reaction.

NOTE: In this Lecture the gas-phase reactions are considered. Aqueous-phase reactions are discussed in Lecture 13.

thermal (or kinetic) reactions : reactions proceeding primarily due to the kinetic energy

of reacting species (i.e., due to thermal agitation of species).

For instance, A + B => products

photochemical reactions : reactions proceeding due to an energy input in the form of

light. For instance, A + hn => products

Figure 6.3 (a, b, c) The fundamental mechanisms of chemical and photochemical processes.

Figure 6.3a. Illustrates a binary, two-bodies, reaction between reactant molecules A and B yielding the product C and D. AB* is an intermediate activated complex.

Figure 6.3b. Illustrates a ternary, three-bodies, reaction between reactant molecules A and B forming AB* a three-way collision with another air molecule, M .

Figure 6.3c. The photodissociation of molecule AB into the products, following the absorption of a photon of radiation.

• Every molecule of a given species has the same potential energy, which is known or readily measurable. In any mixture of different gases, chemical and physical interactions between the gases work to minimize the total potential energy of the system.

Figure 6.4 The potential energy diagram for a binary reaction A+B => C+D

The activation energy is the minimum amount of energy needed for colliding species to react.

The heat of reaction is the potential energy difference between the reactants and products.

• The reaction is said to be exothermic if heat is generated during the reaction. The reaction is said to be endothermic if heat must be absorbed from the environment to make the reaction go.

Most chemical reactions occur by a series of steps called the reaction mechanism.

For example, for the reaction

the mechanism is thought to involve the following steps:

Each of these steps is called an elementary step (or elementary reaction). The substance NO₃ (highly reactive nitrate radical) is known as a reactive intermediate.

• An intermediate is formed in one step and consumed in a subsequent step and so it never seen as a product.

Molecularity is defined as the number of species that must collide to produce the reaction indicated by that step. A reaction involving one molecule is called a unimolecular reaction. Reactions involving the collision of two and three species are termed bimolecular and termolecular, respectively.

Example:

unimolecular reaction: decomposition of peroxyacetyl nitrate (PAN) gives peroxyacyl radical and nitrogen dioxide

CH₃CO(O₂)NO₂ => CH₃CO(O₂) + NO₂

• The rate of a reaction describes how rapidly a chemical change takes place.

Example: consider an elementary reaction A + B => C + D

Rate = - d[A]/dt = - d[B]/dt = d[C]/dt = d[D]/dt

where [ ] means concentration of the species, t is time. Here negative signs denote loss of species A and B.

NOTE: d[A]/dt means {concentration of species A at time t₂ - concentration of species A at time t₁}/ t₂-t₁

The rate of an kinetic or photolytic reaction equals a rate coefficient (or rate constant) multiplied by the concentrations of each of the reactants.

• The rate of photochemical reaction

A + hn => products

is defined as

where J is the photolysis rate coefficient for species A.

NOTE: detailed discussion of the photochemical reactions is given in Lecture 7.

• The rates of unimolecular (one-body) , bimolecular (two-body), and termolecular (three-body) elementary reaction:

unimolecular (one-body): A => products

(reactions caused by thermal decomposition)

bimolecular (two-body): A+ B => products

(any two chemically-active species that collide)

termolecular (three-body): A+ B + C => products (fairly rare reactions)

• In general, the rate of a hypothetical elementary reaction

is defined as

where k is the rate constant (or rate coefficient), [A], [B], and [C] are concentrations of reactants A, B, and C, respectively; and n = a+b+c is the order of reaction.

NOTE: Important to remember that for an elementary reaction, the rate equation follows directly from the molecularity of this reaction (based on the law mass action). Thus a unimolecular reaction is always first order, a bimolecular reaction is always second order, and so on. On the other hand, for an overall reaction containing several elementary steps, experimental determination of the rate is the only means of determining reaction orders.

• Many important atmospheric chemistry reactions are pseudo first order. This means that although they are truly second order reactions, the concentration of one component is essentially constant.

Example: nitric acid formation in air via reaction of nitrogen dioxide with the hydroxyl radical: NO₂ + OH => HNO₃

d[HNO₃]/dt = k₂ [NO₂] [OH] = k* [NO₂] because concentration of OH may be considered constant over some time intervals (hours, but not days). Here k* is called pseudo first order rate constant, and k* = k₂ [OH].

Each kinetic reaction has a rate coefficient associated with it. The rate coefficient is a constant for a given temperature, pressure, and set of reactants, and relates the gas concentrations to the rate of reaction. Such rate coefficients are determined experimentally.

Most reaction rate coefficients are temperature dependent as given by the Arrhenius equation:

where A_r is the collisional prefactor (also known as pre-exponential, or frequency factor, may depend on T), E_r is the activation energy, R is the universal gas constant.

Activation energy is defined as the smallest amount of energy required for reacting species to form an activated complex or transition state before products are formed (see Fig. 6.4 for an illustration).

• concentration of reactants (gases) is in [molec. cm^-3]
• the rate of reaction (regardless of whether it is unimolecular, bimolecular, or termolecular) is in [molec. cm^-3 s^-1]
• the reaction rate coefficients (depend on type of reaction):

for unimolecular reaction is in [s^-1] ;

for bimolecular reaction is in [cm³ molec.^-1 s^-1];

for termolecular reaction is in [cm⁶ molec.^-2 s^-1].