Inorganic Chemistry- I 1. Atomic Theory and Atomic Structure

"You can lead a horse to river, but you can't make him to drink."

Set By: Tessema D. Inorganic Chemistry I module, Wachemo University, Hossana, Ethiopia

Inorganic Chemistry- I

1. Atomic Theory and Atomic Structure

At the end of this chapter, students will be able to:

Have a solidify clue about atomic theories

Have a general overview of quantum mechanics.

Relate the earliest atomic theory with modern atomic theory

Describe the quantum numbers

Familiar with atomic structure and atomic spectra

1.1. Definition Inorganic chemistry is a branch of chemistry that deals with chemical elements and

their compounds excluding hydrocarbons and their derivatives but usually or often

including carbides and other relatively simple carbon compounds especially some

carbon-oxygen and carbon-sulfur compounds (oxides of carbon, metallic carbonates,

and carbon disulfide) and some carbon-nitrogen compounds (hydrogen cyanide and

metallic cyanides). This field covers all chemical compounds except the myriad

organic compounds (carbon based compounds, usually containing C-H bonds), which

are the subjects of organic chemistry. It has enormous applications in every aspect of

the chemical industry including catalysis, materials science, pigments, surfactants,

coatings, medicine, fuel, and agriculture. The inorganic realm is extremely board

which providing essentially limitless areas for investigation. The inorganic chemistry

has relation with other chemistry courses. For instance, Organ metallic chemistry: a

bridges organic and inorganic chemistry. It bridges organic compounds bridges with

inorganic compounds that contains direct M-C bond. Bioinorganic chemistry bridges

biochemistry with inorganic chemistry. Environmental chemistry studies both organic

and inorganic compounds.

Elementary substances and solid-state inorganic compounds are widely used in the

core of information, communication, automotive, aviation and space industries as well

as in traditional uses. These inorganic compounds are also indispensable in the

frontier chemistry of organic synthesis using metal complexes, homogeneous catalysis,

bioinorganic functions, etc. The distinction between the two disciplines is far from

absolute, and there is much overlap, most importantly in the sub-discipline of

organometallic chemistry.

http://en.wikipedia.org/wiki/Chemical_compound

http://en.wikipedia.org/wiki/Organic_compound

http://en.wikipedia.org/wiki/Organic_chemistry

http://en.wikipedia.org/wiki/Organometallic_chemistry



1.2. Introduction to Atomic-Theory

Atomic theory describes about the historical models of the atom, especially the

external structure of atoms and how atoms combine to form molecules. This theory is

a scientific theory that deals the nature of matter; states that matter is composed of

discrete units which called atoms. Atom is the smallest particle into which an element

can be divided and still be the same substance. Atoms are so small until recently, no

one had ever seen it. But ideas, or theories, about atoms have been around for over

2,000 years. The theory of the atom has had a long history. The ancient Greeks

postulated that matter exists in the form of atoms. But they did not base their theory

on experiment and they cannot develop additional ideas about atoms. Different

scholars (researchers) give their idea on the atomic theory.

In order to understand the structure of an atom, different types of atomic models were

developed using different instruments. Some of the models are:

1. Democritus theory (300 BC)

The Greek philosopher Democritus expressed his own postulate and he states that:

Matter consists of very small, indivisible particles which are called as atoms”.

Atoms are derived from Greek words ATOMOS which mean indivisible into smaller

particle. Atoms are indivisible particles. Explains certain natural occurrences such

as the existence of elements

Limitation of Democritus atomic theory

His theory does not support by experiment but only theoretical view

Does not have any information about subatomic particles (electron, proton and neutron).

Due to the lack of experiment, Democritus’ idea was not accepted by many of researchers or

scholars because Experimental evidence from early scientific investigations provided

support for the notion of ―atomism‖ and gradually gave rise to the modern definitions of

elements and compounds.

2. Dalton’s Atomic Theory

John Dalton (1766-1844), an English schoolteacher, developed the first useful atomic

theory of matter around 1808. His findings were based on experiments and also from

laws of chemical combination. He used fundamental laws of chemical combination

just described as the basis of an atomic theory. His theory involved many

assumptions:

All substances are composed of tiny, indivisible particles which called ―atoms‖.



Atoms of the same element are identical in size, mass and properties but, atoms of

different elements have different properties (he doesn’t recognize isotopes).

Chemical reactions only rearrange the way that atoms are combined; the atoms

themselves are unchanged. Dalton realized that atoms must be chemically

indestructible for the law of mass conservation to be valid. If the same numbers and

kinds of atoms are present in both reactants and products, then the masses of

reactants and products must also be the same. This means that: Compounds are

formed by the union of two or more different elements

Dalton model is known as the solid sphere model (like billiard balls).

Importance and Improvement on previous model

Explains how atoms combine to form molecules.

Explains chemical change better than the particle theory.

Define conservation of mass and definite proportion.

Dalton’s theory put down a corner stone for modern atomic theories since he uses the

law of chemical combination.

Laws of Chemical Combination

Various chemical reactions take place according to certain laws, known as the Laws

of chemical combination. There are three common laws of chemical combinations.

These are:

(I) Law of conservation of mass/the law of mass action

It also known as the law of indestructibility of matter since this law states that

―matter is neither created nor destroyed in the course of chemical reaction rather it

may change from one form to other‖. The total mass of materials after a chemical

reaction is same as the total mass before reaction. In short this law states that

―during any physical or chemical changes, the total mass of product equal to the

total mass of reactants‖. For example, in an experiment 63.5g of copper combines

with 16g of oxygen to give 79.5g of cupric oxide (a black oxide of copper).

(II) Law of constant or definite composition: According to this law, in a given

compound, the constituent elements are always combined in the same proportions by

mass, regardless of the origin or mode of preparation of the compound. What this law

means is that when elements react chemically, they combine in specific proportions,



not in random proportions. A sample of pure water which obtains river, sea, well,

spring whatever the source, always contains 88.9% by mass of oxygen and 11.1% by

mass of hydrogen.

(III) Law of multiple proportions: According to this law, when two elements A and B

combine to form more than one chemical compounds then different weights of A,

which combine with a fixed weight of B, are in proportion of simple whole numbers. Or

when two/more elements combine to form two/more compounds, the mass of one

element combines with a fixed mass of another element. Example: Carbon monoxide

(CO): 12 parts by mass of carbon combines with 16 parts by mass of oxygen. Carbon

dioxide (CO2): 12 parts by mass of carbon combines with 32 parts by mass of oxygen.

Ratio of the masses of oxygen that combines with a fixed mass of carbon (12 parts) is

16: 32 or 1: 2

Another example, nitrogen and oxygen can combine either in a 7:8 mass ratio to

make a substance we know today as NO, or in a 7:16 mass ratio to make a substance

we know asNO2. Clearly, the second substance contains exactly twice as much oxygen

as the first.

NO : 7 g nitrogen per 8 g oxygen N:O mass ratio = 7:8

NO2 : 7 g nitrogen per 16 g oxygen N:O mass ratio = 7:16

This result makes sense only if we assume that matter is composed of discrete atoms,

which combine with one another in specific and well-defined ways.

Limitations of Dalton’s atomic theory

Does not include the existence of the nucleus

Does not explain the existence of ions or isotopes

Does not have any information on subatomic particles (electron, proton and neutron).

3. J. J. Thomson theory (1850)

This model is known as raisin model or the chocolate chip cookie model since the

electrons are embedded on the solid positive nucleus. He proposed that the atom is a

solid sphere of positively charged material and the negatively charged electrons spread

throughout the atom which similar to plums in a pudding or chocolate chips in ice

cream. Atoms are solid spheres made-up of a solid positive mass or core. It also

contains tiny negative particles embedded in the positive core.



He was the first scholar to discover the presence of electron in atom of an element by

the help of cathode ray tube experiment.

Limitation

Does not explain the where the electrons are found (energy levels).

Does not explain the role of electrons in bonding (valence electrons)

Does not explain about neutrons (radioactivity and isotopes)

4. Rutherford Atomic Theory (1905)

This model is known as the planetary model, since the electrons revolve around the

nucleus just like the planets do around the sun. According to Famous gold leaf

experiment, the nucleus is positively charged and the (electrons) are revolving outside

the nucleus.

Importance and Improvement on previous model

First real modern view of the atom.

Explains why the electron spins around the nucleus

Proposes that the atom is really mostly empty space

Rutherford determines true nature of atom and the phenomenon of radioactivity

(decay of unstable atomic nuclei)

Limitation

Does not place electrons in definite energy levels around the nucleus.

Doesn't include neutrons in the nucleus.

Does not relate the valence electrons with atomic charge.

5. Bohr-theory (1920)

He states that electrons are found in definite energy levels around the nucleus. Used

atomic spectra to prove that electrons are placed in definite orbital’s (called shells)

around the nucleus.



Several questions are arisen on Bohr atomic theory like:

1. When electrons are revolving around the nucleus as a planet in planetary motion they might

be slow down gradually to the nucleus or escape from the nucleus? He answered as: An

electron did not radiate energy if it stayed in one orbit, and therefore did not slow

down.

2. Why should electrons move in an orbit round the nucleus? He replied as: When an electron

moved from one orbit to another it either radiated or absorbed energy. If it moved towards

the nucleus energy was radiated and if it moved away from the nucleus energy was

absorbed.

3. Since the nucleus and electrons have opposite charges, they should attract each

other. He answered as: an electron to remain in its orbit the electrostatic attraction

between the electron and which tends to pull the electrons towards nucleus must

be equal to the centrifugal force.

Bohr's basic theory deal that electrons in atoms can only be at certain energy levels,

and they can give off or absorb radiation only when they jump from one level to

another. In his model, an atom consists of an extremely dense nucleus (p and n) and

surrounded by electrons that travel in set orbits around the nucleus. He hypothesized

that the energy possessed by these electrons and the radius of the orbits are

quantized, meaning it is limited to specific values and is never found between those

values. These ―orbits‖ have energies that dependent on their distance from the

nucleus. Mathematically, this model was determined by the equation:

)11

(22

fi

Hnn

RE , where RH is the Rydberg constant (2.18 x 10-18J), ni and nf are

integers with nf > ni.



From mathematical equations describing the orbits for the hydrogen atom, together

with the assumption of quantization of energy, Bohr was able to determine two

significant aspects of each allowed orbit:

1. Where the electron can be found with respect to the nucleus that is, the radius, r, of

the circular orbit is given by: r=aon2 where n is a positive integer (1, 2, 3...) that tells

which orbit is being described and a0 is the Bohr radius. Bohr was able to calculate

the value of a0 from a combination of Planck’s constant, the charge of the electron,

and the mass of the electron as a0=5.292*10 -11 m =0.5292 Å.

2. The stable electron in its orbit has a potential energy, E. This is given by:

Here h=Planck’s constant, m=the mass of the electron, and the other symbols have the

same meaning as before. E is always negative when the electron is in the atom; E=0

when the electron is completely removed from the atom (n=infinity).

Different regions of hydrogen spectrum

Bohr’s theory did make important contributions;

1. It suggested a reasonable explanation for the discrete line spectrum of the elements

electron

nucleus

Allowed Bohr orbital’s of hydrogen atom



3. It introduced the idea of quantized electron energy levels (orbits).

4. Explains the role of valence electrons in bonding

5. Places electrons in definite energy levels as 2 electrons in the first, 8 electrons in

the second etc.

Limitation:

It does not explain the shapes of molecules or other abnormalities that result in the

formation of unevenly shared pairs of electrons. Example: The difference carbon-

carbon bonds between diamond and graphite.

Bohr’s theory only works for hydrogen atoms and if there are more than one

electron, the calculations for the electron energy and the orbit radii breakdown.

Therefore, a new theory needed for multi-electron elements.

1.3. Modern Atomic Theory This model is known as Quantum Mechanical Model or Electron Cloud Model. It

states that ―electrons are always moving around the nucleus in a "cloud" of energy

levels”. It states that the small particle atom is consists of three parts such as,

proton, neutron and electron. The electrons are found around the nucleus whereas;

protons and neutrons are found in the central part of atom which is called nucleus.

As a result, the modern theory describes that, the mass of one atom is the summation

of the number of proton and the number of neutrons. NB: the mass of an atom is

occupying by the nucleus whereas the large space of the atom is occupying by

electrons.

Quantum Model of Atom

Bohr's theory worked well for hydrogen atom (one electron), but the theory is failed

atoms with more than one electrons. The developing experimental science of atomic

spectroscopy provided extensive data for testing of the Bohr atomic theory and its

modifications that forced the theorists to work hard to explain the spectroscopists’

observations. In spite of their efforts, the Bohr atomic theory eventually proved

unsatisfactory; the energy levels are valid only for the hydrogen atom. So that an

important characteristic of the electron, its wave nature, still needed to be considered.

According to the de Broglie equation, proposed in the 1920s, all moving particles

have wave properties described by the equation; mv

h where λ=wavelength of the

particle, h = Planck's constant, m = mass of the particle, v = velocity of the particle.



Particles massive enough to be visible have very short wavelengths, too small to be

measured. Electrons, on the other hand, have wave properties because of their very

small mass.

Electrons moving in circles around the nucleus (in Bohr’s Atomic theory), can be

thought of as forming standing waves that can be described by the de Broglie

equation. However, we no longer believe that it is possible to describe the motion of an

electron in an atom so precisely. This is a consequence of another fundamental

principle of modern physics, Heisenberg's uncertainty principle, which states that

there is a relationship between the inherent uncertainties in the location and

momentum of an electron moving in the x direction:

4

hpx x , where, Δx= uncertainty in the position of the electron, Δp= uncertainty

in the momentum of the electron. The energy of spectral lines can be measured with

great precisions which in turn allowing precise determination of the energy of

electrons in atoms. This precision in energy also implies precision in momentum (Δp,

is small); therefore, according to Heisenberg, there is a large uncertainty in the

location of the electron (Δx is large). These concepts mean that we cannot treat

electrons as simple particles with their motion described precisely, but we must

instead consider the wave properties of electrons, characterized by a degree of

uncertainty in their location. In other words, instead of being able to describe precise

orbits of electrons, as in the Bohr Theory, we can only describe orbitals, regions that

describe the probable location of electrons or the region of space where the electrons of

an atom or molecule are found.

In the 1920’s Erwin Schrödinger applied the principles of wave mechanics to atoms

and developed the Quantum Mechanical Model of the Atom. Basically, Schrödinger

said to give up on the idea of literal orbits for the electrons and instead concentrate on

the electron as a wave. This theory builds on Bohr’s idea of quantized energy levels

(orbits) and adds additional requirements for electron location and energy. Working

with Heisenberg’s Principle, Schrödinger developed a compromise which calculates

both the energy of an electron and the probability of finding an electron at any point in

the molecule. This is accomplished by solving the Schrödinger equation, resulting in

the wave function, Ψ.



These regions were termed as orbitals. The Schrodinger equation describes the wave

properties of an electron in terms of its position, mass, total energy, and potential

energy. The equation is based on the wave function, Ψ, which describes an electron

wave in space; in other words, it describes an atomic orbital. In its simplest notation,

the equation is: HΨ=EΨ, where H= the Hamiltonian operator, E= energy of the

electron, Ψ= the wave function. This concept is based on the quantum numbers.

Schrodinger wave equation

Schrödinger wave equation of an electron wave moving in any of three dimensions

(axes) x, y, z its wave motion can be described by the wave equation:

*

(

) +

OR

Where, h = Planck's constant, m = mass of the particle (electron), e = charge of the

electron, x, y, z = are Cartesian coordinates distance from the nucleus and Z = charge

of the nucleus

√

(

)

√

{

}

{

}

Where,√

The potential energy V is a result of electrostatic attraction between the electron and

the nucleus. Attractive forces, like those between a positive nucleus and a negative



electron, are defined by convention to have a negative potential energy. An electron

near the nucleus (small r) is strongly attracted to the nucleus and has a large negative

potential energy. Electrons farther from the nucleus have potential energies that are

small and negative. For an electron at infinite distance from the nucleus , the

attraction between the nucleus and the electron is zero, and the potential energy is

zero. Because every matches an atomic orbital, there is no limit to the number of

solutions of the Schrodinger equation for an atom. Each describes the wave

properties of a given electron in a particular orbital. The probability of finding an

electron at a given point in space is proportional to 2.

Unique characteristics of wave function are given below.

S.

No

Symbol of Wave Function Definition

1 The wave function Ψ must

be single valued.

There cannot be two probabilities for an electron at any

position in space.

2 The wave function Ψ and

its first derivatives must be

continuous.

The probability must be defined at all positions in space

and cannot change abruptly from one point to the next.

3 The wave function Ψ must

approach zero as r

approaches infinity.

For large distances from the nucleus, the probability

must grow smaller and smaller (the atom must be

finite).

4 The integral

∫ Ψ Ψ

The total probability of an electron being somewhere in

space = 1. This is called normalizing the wave function.

5 The integral

∫ Ψ Ψ

All orbitals in an atom must be orthogonal to each

other. In some cases, this means that the orbitals must

be perpendicular, as with the px, py, and pz orbitals.

1.4. Quantum Numbers The quantum numbers are parameters that describe the distribution of electrons in

the atom, and its fundamental nature such as the size, energy, shape and position of

electrons. There are four fundamental Quantum Numbers. These are:

1. Principal Quantum Number (n): Represents the main energy level, or shell,

occupied by an electron. It is always a positive integer that is n= 1, 2, 3...∞



2. Angular Momentum Quantum Number (l): Represents the energy of sublevel, or

type and shape of orbital occupied by the electron. The value of l depends on the value

of n such that l= 0, 1,... n-1.

L value 0 1 2 3

Type S P D f

Shape Spherical Dumble bell Double dumble bell complicated

3. Magnetic Quantum Number (ml): Represents the number of possible orientations

in 3-D space for each type of orbital. It determines the number of orbitals and has the

value of ml ranges between –l and +l including zero such that ml= -l ...0 ...+l.

4. Spin Quantum Number (mS): Represents the two possible orientations that an

electron can have in the presence of a magnetic field, or in relation to another electron

occupying the same orbital. Only two electrons can occupy the same orbital, and they

must have opposite spins. When this happens, the electrons are said to be paired. The

allowed values for the spin quantum number ms are +1/2 and -1/2. It identifies the

type of rotations (clockwise/upward or anticlockwise/downward) direction.

1.5. Electronic Configuration

Electron configuration is the way of filling/distribution of electrons among the given

orbitals. In writing electronic configurations or electron filling of orbitals, we follow the

Aufbau principle, Hund’s rule, and the Pauli Exclusion Principle. These are:

A. The Aufbau Principle (building up principle): explains the order of filling

electrons in various orbitals of an atom. Filling begins with the orbitals in the

lowest energy, or most stable, shells and continues through the higher-energy

shells, until the appropriate number of orbitals is filled for each atom. Thus, the 1s

orbital fills first, then the 2s, followed by the 2p and the 3s orbitals.



B. Hund's rule: states that each degenerate orbital, (e.g.2px, 2py, and 2pz) must first

receive one electron before any of the orbitals can receive a second electron.

C. Pauli Exclusion Principle: states two important statements.

Each orbital contains a maximum of two electrons. These two electrons must have

opposite values for the spin, which is generally indicated by showing the electrons

as arrows pointing up () or down ().

Not two electrons can have the same four quantum numbers; at least they differ

by one quantum number. Example: 2He: 1s2; n=1, l=0, ml=0, but they differ by

spin quantum number. One is upward while the other is downward in position.

D. Made lung's Rule: This rule 'explains' why the 4s orbital has a lower energy than

the 3d orbital. Orbital energy is sometimes dependant both principal quantum

number and l is the subsidiary quantum number; so orbitals fill with n + l energy

sequences.

Example 2; Nitrogen (N) has z=7, its E.C is 1s2, 2s2, 2p3, the orbitals are filled as:

Potential Energy

1s2s

2px 2py 2pz

N (z=7) 1s2, 2s2, 2px1 2py

1 2pz1

Example 3; Oxygen (O) has z=8, its electronic configuration is 1s2, 2s2, 2p4, the orbitals

are filled as follows.

Potential Energy

1s2s

2px 2py 2pz

O (z=8) 1s2, 2s2, 2px2 2py

1 2pz1

Remarks:

1. The maximum number of orbitals for any principal quantum number (n) is n2.

2. The maximum number of electrons for any principal quantum number (n) is 2n2.

3. The maximum number of orbitals for any angular momentum quantum number (l) is 2l+1.



4. The maximum number of electrons for any angular momentum quantum number (l) is 4l+2.

Exercise

1. Write the electronic configuration and fill the orbitals of the following elements, He, B, F, Ne,

Cr, Cu and S.

2. Write the possible four quantum numbers for 4d orbitals.

3. How many maximum electrons can hold an orbitals with quantum numbers:

A. n=4, l= 2,ml=1,ms=+1/2

B. n=4, l= 2,ml=1

C. n=4, l= 2

D. n=4

1.6. Atomic Structure By combining different atomic theories which have been discussed in previous section,

we can deduce that any atom consists of a central denser nucleus (proton and

neutron) and the surrounding electrons around the nucleus.

electrons

nucleus

The nucleus is made up of the neutrons (carry zero charges) and the protons (carry

positive charges). Therefore, the nucleus is positively charged particle.

All atoms are made from three subatomic particles (Protons, neutron and electrons).

The three fundamental particles have the following properties.

Particle Charge Mass (g) Mass (amu)

Proton +1 1.6727 x 10-24 g 1.007316

Neutron 0 1.6750 x 10-24 g 1.008701

Electron -1 9.110 x 10-28 g 0.000549

Protons and neutrons have almost the same mass, while the electron is approximately

2000 times lighter. Protons and electrons carry charges of equal magnitude, but



opposite charge. Neutrons carry no charge (they are neutral). The positive charges on

the protons are always balanced by the negative charges on Electron in neutral atoms.

Electrons move rapidly around the nucleus and constitute almost the entire volume of

the atom.

Chemical reactions involve either the transfer or the sharing of electrons between

atoms. Therefore, the chemical reactivity/ properties of an element are primarily

dependent upon the number of electrons in an atom of that element.

Therefore, we can say that the chemical reactivity of an atom is dependent upon the

number of electrons (valence electrons).

Atomic structure refers to the configuration of the constituents of the atom. When we

speak of electronic configuration we deal with information that tells us where in the

atom the electron(s) is (are) found. The chemistry of an element is determined largely

by its electronic configuration, so it is very important to get as much information as

possible about the behavior of electrons in atoms. It is a good idea to start with atom,

the smallest constituent particle of matter, which consists of dense central nucleus

surrounded by a cloud of negatively charged electrons. Atoms are composed of smaller

subatomic particles in an atom. These are: proton, neutron, and electron. The

number of protons in an atom identifies which element that atom is and gives that

element its atomic number. The number of protons in the nucleus and the

corresponding number of electrons around the nucleus controls each element's

chemical properties. However, the electrons are the active portion of an atom when it

chemically bonds with another atom. The electrons determine the structure of the

newly formed molecule. The number of protons in the atoms that make up a sample of

a particular element is always the same, but the number of neutrons can vary.

Example; Oxygen, has 16 atomic mass (written below the symbol of the atom) and 8

atomic number (Z) (written on the upper left the symbol of the atom) and contains 8

neutrons. This can be represented as follows.

Mass Number: Mass number is the number of protons and neutrons in the nucleus of

an atom. The mass of an atom essentially depends upon the number of protons and

neutrons in the nucleus.

Mass number = number of neutron +number of proton/electron



N.B the number of electrons exist around the nucleus and the number of protons

which are found in the central part of atom is equal if the element is neutral.

A complete symbolism or shorthand for describing atoms which is universally used

across all scientific disciplines can be denoted as;

Mass number A

X Chemical symbol

Atomic number Z

Example: If an element has an atomic number of 34 and a mass number of 79, what

is the: a) number of protons , b) number of neutrons , c) number of electrons and

d)represent the complete symbol of the element.

Solution: atomic number = 34g

Mass number = 79g/mol therefore, from the definition of atomic number, the number

of protons in the nucleus is the same with the atomic number of atom.

Number of proton = atomic number =number of electron

a) Number of proton =34g

b) Number of neutron = mass number – number of protons

79g/mol – 34 = 45 neutrons

c) Number of electron =34g

d) So the symbol the element is represented by Au79

34

Isotopes: Dalton was wrongly deduced about all elements of the same type being

identical. Atoms of the same element can have different numbers of neutrons but the

same number of protons is called isotopes. Frederick Soddy (1877-1956) proposed the

idea of isotopes in 1912 and he states that ―Atoms with the same number of protons

but different numbers of neutrons or Isotopes are atoms of the same element having

different masses, due to varying numbers of neutrons.

Example: Hydrogen contains three isotopes

1. protium: contains one proton and one electron

2. deuterium :one proton, one electron and one neutron

3. Tritium: consists of one proton, one electron and two neutrons.

Average atomic Mass: The average atomic mass is based on the abundance

(percentage) of each variety of that element in nature. To calculate the average atomic



mass multiply the atomic mass of each isotope by its Abundance (expressed as a

decimal), then add the results.

Average atomic mass = %abundance/100 X accurate atomic mass

Example: Calculate the value of Average atomic mass for naturally occurring chlorine

if the distribution of isotopes is 75.77% 3517Cl and 24.23% 37

17Cl. Accurate masses for

35Cl and 37Cl are 34.97 and 36.97, respectively.

The relative atomic mass of chlorine is the weighted mean of the mass numbers of the

two isotopes:

Average atomic mass = (75.77/100X34.97) + (24.23/100X36.97)=35.45

Unit Summary Atomic theory is developed by different scholars: Democritus (atoms are indivisible);

Dalton’s atomic theory states three ideas: I) matter is composed of indestructible

particles called atoms; II) atoms of an element are identical to one another but

different from atoms of all other elements, and III) chemical compounds are

combinations of atoms of different elements. Based on this theory, Dalton proposed

still another law of chemical combination, the law of multiple proportions. J.J.

Thomson discovered the present of one -ve particles in an atom using cathode ray

experiment later on Rutherford discovered the present of denser nucleus in an Aton

using gold leaf experiment and develop planetary motion. Borhr,s also modified the

planetary model using emission and absorption of atoms as: Energy is emitted or

absorbed by an atom only when an electron moves from one orbit or energy level to

another unless otherwise an electro doesn’t emit or absorb energy if it stays in one

orbit. This leads the quantization of energy, orbitals or radius of orbitals. Quantum

mechanical model develops a set of scientific principles describing the known behavior

of energy and matter that predominate at the atomic and subatomic scales by different

scholars such as deBorglie (dual nature of electrons), Shrodinger (wave function),

Heinserberg (uncertainty principle or errors in measurements) and other scholars.

This electron cloud model leads the four quantum numbers specify completely the

position of electron in a given atom. They give its position in the main energy level (n),

the sub-energy level (l), the orientation of its orbital (ml) and the direction of its spin

(ms). In other words, they serve as an address for the electron as electron clouds. By

using these quantum numbers and some regulations like Pauli principle (the

maximum electros in one shell should be two, not two electrons can have the same



four quantum numbers); Aufbau principle (energy sequences) and Hud’s rule

(degenerate orbitals should receive the 1st electron before any of them can received the

2nd electros), we can distribute electros in their house (orbitals).

Review Questions 1. When an iron object rusts, its mass increases. When a match burns, its mass decreases.

Do these observations violate the law of conservation of mass? Explain.

2. What is basic difference between wave function of electron and square of wave function

( of electron ( 2.

3. It is possible to determine the exact location of an electron? Explain.

4. What evidence supports the idea that electrons are particle and wave natures?

5. What are the values of n, l and ml for the following sub shells? 1s, 4s, 3p, 3d, and 4f?

6. State the Pauli Exclusion Principle. Would any of the following electron configurations

violate this rule? 1s2,1s22p7, 1s3 explain.

7. State Hund’s rule. Would any of the following electron configurations violate this rule?

1s2, 1s22s22px2 ,1s22s22px

12py1, 1s22s22px

12pz1, 1s22s22px

12py12pz

1

8. What are the maximum orbital and electrons for the following quantum numbers?

I. n=3, l=2, ml=0,1

II. n=3, =0,1

III. n=3, l=2, ml=0,1, ms=±1/2

2. Chemical Bonding

At the end of this chapter, students will be able to:

Explain the basic concepts of chemical bonding

Discuss on type of bonding, bond formation and geometry of molecules.

Familiar with different bonding theories (Lewis, VSEPRT,VBT, hybridization, MOT)

Familiar with solid structure and crystal defect

Familiar with Bond theory and band energy

Brainstorming Questions

1. Describe why and how chemical bonding is formed?

2. Discus about type of chemical bonds and Define chemical bond.

3. Why we want to sit rather than standing? Why we want to sleep rather than to sit?

Why we want to emigrate abroad like America and Europe? Take time for

discussion these events with relation of chemical bond.



2.1. Chemical Bonding

Bonding is the joining of at least two atoms to form a molecule or compound. The

electrons in the valence shell are the active portion of an atom during bonding. A

chemical bond is a mutual electrical attraction between the nuclei and valence

electrons of different atoms that binds the atoms together.

Why atoms combined to form compounds, clusters, molecules/ polyatomic ions?

to attain octet (eight) rules

to have lower energy (to be stable)

to have net attractive force

e

ee

e

+ +

R

R

R

A A

AA

where, A= attraction force, R= repulsion force

Chemical bond is formed when A>R:

Example: H + H H2(g) but, He + He no reaction since R>A.

Most atoms are chemically bonded to each other? As independent particles, they are at

relatively high potential energy. However, nature favors arrangements in which

potential energy is minimized. This means that most atoms are less stable existing by

themselves than when they are combined. By bonding with each other, atoms

decrease in potential energy, thereby creating more stable arrangements of matter.

How do atoms combined to form compounds/ polyatomic ions?

In chemical bonds, atoms are combined to form compounds, clusters, molecules/

polyatomic ions by:

Transferring /losing of their valence electrons.

Gaining of extra valence electrons

Sharing of valence electrons

Coordinating valence electrons

2.2. Types of Chemical Bonding When atoms bond, their valence electrons are redistributed in ways that make the

atoms more stable. The way in which the electrons are redistributed determines the



type of bonding. There are different types of chemical bonds that results by transfer or

sharing of electrons. These are:

Electropositive element +Electronegative element

Electronegative elment +Electronegative element

Electropositive element +Electropositive element

Ionic bond

Covalent bond

Metallic bond

2.2.1. Ionic Bond/Electrovalence bond Ionic bond is a type of chemical bond that results from the electrostatic force of

attraction between two oppositely charged ions. It is formed due to transfer of

electrons from one atom (mostly metals) to another (mostly nonmetals) by losing and

gaining of electrons. The atom that loses electrons will form a cation and the atom that

gains electrons will form an anion. These oppositely charged ions come closer to each

other due to electrostatic force of attraction and thus form an ionic bond. In general,

an ionic bond is formed between two atoms that contain greater electronegativity

difference and it is between a metal atom and a nonmetal atom. Example: NaCl, LiF,

MgCl2.

Favorable condition for ionic bond formation

Low ionization energy of the metals

High electro-affinity of the nonmetals

High lattice energy

High electro-negativity difference between a metal and a nonmetal atom

HABER –BORN CYCLE: The Born-Haber cycle allows us to understand and determine

the lattice energies of ionic solids. Lattice Energy is a type of potential energy that may



be defined in two ways. In one definition, the lattice energy is the energy required to

break apart an ionic solid and convert its component atoms into gaseous ions. This

definition causes the value for the lattice energy to always be positive, since this will

always be an endothermic reaction. The other definition says that lattice energy is the

reverse process, in which energy is released when gaseous ions bind to form an ionic

solid. In this definition, lattice energy will be negative and the process is always

exothermic. Its values are usually expressed with the unit of kJ/mol. There are several

important concepts to understand before the Born-Haber Cycle can be applied to

determine the lattice energy of an ionic solid; ionization energy, electron affinity,

dissociation energy, sublimation energy, heat of formation, and Hess' Law.

Ionization Energy (IE) is the energy required to remove an electron from a gaseous

atom or ion. This process always requires an input of energy, and thus will always

have a positive value. It is the measure of how much the outer most electrons held by

nuclear attraction. M (g) +IE → M+ (g) +e−

Electron Affinity (EA) is the energy released when an electron is added to a gaseous

atom or ion. Since energy is released its value is negative value.

X(g)+ e X-(g) + EA

Heat of dissociation energy (∆Hdiss) is the energy required to break apart a

compound/ molecules into elemental state. It is also known as heat of

demolecularazation: molecule X2(g) ∆Hdiss X (g)

The dissociation of a compound requires an input energy and is always an

endothermic process (positive).

Sublimation energy (∆Hsub) is the energy required to change of phase from solid to

gas, by passing the liquid phase. M (s) ∆Hsub M(g)

This is an input of energy, and thus has a positive value. It may also be referred to as

the energy of atomization.

The heat of formation (∆Hrxn) is the change in energy when a compound forming

from the given elements. This may be positive or negative, depending on the atoms

involved and how they interact. If it is negative the reaction is exothermic while it is

endothermic if it is positive.

UEAIEHHH subdissrxn

http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Electron_Affinity

http://chemwiki.ucdavis.edu/Physical_Chemistry/Thermodynamics/State_Functions/Enthalpy/Heat_of_Sublimation



Example: The lattice enthalpy change (U) in the formation of an ionic lattice from the

gaseous isolated sodium and chloride ions is -788 kJ/mole. This enthalpy change,

which corresponds to the reaction Na+(g) + Cl-(g) NaCl(s), is called the lattice energy

of the ionic crystal. For sodium chloride the heat of formation and the Born - Haber

cycle is:

2.2.2. Covalent Bond Covalent is the chemical bond formed between two atoms due to the sharing of electron

pair(s). It is formed between two atoms that contain the same or small difference in

electronegativity. Bonding between two atoms of the same element is completely covalent.

In covalent bond, single, double, triple and quadruple bonds are formed. There are

different types of covalent bonds like polar and non-polar and coordinate covalent bonds.

A polar-covalent bond is a covalent bond in which the bonded atoms have an

unequal attraction for the shared electrons. Example: hydrogen-chlorine bond (H+-

Clδ⁻). Polar covalent bonding occurs because one atom has a stronger affinity for



electrons than the other. When electrons are shared but shared

unequally. Non-polar covalent bond is a covalent bond in which the bonding

electrons are shared equally by the bonded atoms, resulting in a balanced distribution

of electrical charge. Example; hydrogen-hydrogen bond (H: H), Cl2. The electrons are

shared equally.

Coordinate Covalent bond: A coordinate covalent bond (also called a dative bond) is

formed when one atom donates both of the electrons to form a single covalent bond.

These electrons originate from the donor atom as an unshared pair. Example:

BF

F

F

+ N H

H

H

B N

F

F

F

H

H

H

Dative bond

2.2.3. Metallic Bond Brainstorming Questions: why metals are conductors and non-metals are non-

conductors of heat and electricity?

It is the attraction between metal atoms in a metallic crystal. It is formed between

electropositive metal atoms of same or different elements. The valence electrons of

pure metals are not strongly associated with particular atoms. This is a function of

their low ionization energy. Electrons in metals are said to be delocalized (not found

in one specific region, such as between two particular atoms). Since they are not

confined to a specific area, electrons act like a flowing ―sea‖, moving about the

positively charged cores of the metal atoms. Delocalization can be used to explain

conductivity, malleability, and ductility. In general, the greater the number of

electrons per atom that participate in metallic bond the stronger the metallic bond.

E.g. the metal atoms Na, Cu, Ag, Fe etc. are bound to each other in their crystals by

metallic bonds.



2.3. Comparison of Ionic versus Covalent Bonding Formed between high electropositive

metal and high electronegative nonmetals

They are solids with high melting point

(typically _400°C). Soluble in polar

solvents such as water, but insoluble in

nonpolar solvents( C6H14, CCl4)

Molten /aqueous compounds conduct

electricity well because they contain

mobile charged particles (ions).

They are often formed between two

elements with quite different EN,

usually a metal and nonmetals

Formed between two electronegative

nonmetals

They are gases, liquids, or solids with low

melting points (typically=300°C)

Many are insoluble in polar solvents such as

water. Most are soluble in nonpolar solvents,

such as hexane, and CCl4.

Molten compounds don’t conduct

electricity because they don’t contain

mobile charged particles (ions).

They are often formed between two elements

with similar (EN), usually nonmetals

2.3.1. Polarizing power and polarizability (Fajans’ Rule) Ionic and covalent bonding are the two extreme types of chemical bonding, and almost

the other remain bonds are formed intermediate in type. This is explained in terms of

polarizing (that is deforming) the shape of the ions.

The type of bond between A+ and B- depends on the effect one ion has on the other.

The positive ion attracts the electrons on the negative ion and at the same time it

repels the nucleus, thus distorting or polarizing the negative ion. The negative ion also

polarize the positive ion, but since anions are usually larger, and cations small, the

effect of the large ion on a small one will be much less pronounced. If the degree of

polarization is quite small, then the bond remains largely ionic. If the degree of

polarization is large, electrons are drowning from the negative ion towards the positive

ions, resulting in a high concentration of electrons between the two nuclei, and a large

degree of covalent character result.

Fajans put forward four rules which summarize the factors favoring polarization and

covalence (covalent compounds).

A small positive ions favors covalence: In small ions the positive charge is

concentrated over a small area. This makes the ion highly polarizing power, and very

good at distorting the negative ions. Example: LiCl <BeCl2, <BCl3<CCl4 in covalent

nature since the size of cations decrease from Li+ to C+4



2. A large negative ion favours covalence: Large ions are highly polarizable, that is

easily distorted by the positive ion, because the outermost electrons are shielded from

the charge on the nucleus by filled shell of electrons. Example: HF<HCl< HBr<HI in

covalent character since the polarizability and the size of anions increased as the order

of F-< Cl-< Br-<I-.

3. Large charge on either ion or on both ions favors covalence: This is because a

high charge increases the amount of polarization.

Example: CH4>NH3>H2O>HF in covalent character since the polarizability and the

charges anions decreased as the order of C4->N3->O2-> F-.

4. Polarization, and hence covalence: is favored if the positive ion does not have a

noble gas configuration. A noble gas configuration is the most effective at shielding

the nuclear charge, so ions without the noble gas configuration will have high charges

at their surface, and thus be highly polarizing.

Generally, the polarizing power of cation increases as the ions becomes smaller and

more highly charged. The polarizability of a negative ion is greater than that of the

positive ion since the electrons are less firmly bound because of the differences in

effective nuclear charge. Large negative ions are more polarizable than the small once.

Polarizing power: the ability of cation to distorted/ withdraw electron from the

anions. Polarizing power increase across the period since the size of cation decreased

and the positive charges of cation increased. Example: NaCl< MgCl2< AlCl3 in covalent

nature since the size of cation decreased as Al3+,Mg2+,Na+ ,but the charge increase from

Na to Al. Polarizing power decreased down the group since the size of cation

increased. Example: NaCl> KCl> CsCl in polarizing power since the size of cation

increased from Na to Cs down the group.

Polarizability: the ability of anion to distort/ deform by caions. Polarizability of

anion decrease across the period since the size of anion decreased and the negative

charges of anion decreased. Polarizability of anion increased down the group since

the size of anion increased.



Metallic bond

Covalent bond

favord by: high positive charge,small cation size and large anionsize

Ionic bond

Favord by: low +ve charge,Large cation size and Small anion size

Application of Fajan,s rule: Fajan,s rule use to distinguish the ionic and covalent

character of a given compounds based on the size of cation and anion, charge of both

cation and anion. It also used to differentiate the solubility of compounds based of

―like dissolve in like principle‖.

2.4. Shapes of Simple Covalent Molecules Simple covalent molecules are molecules that are held together by relatively weak

intermolecular forces of attraction, although the covalent bonds within the molecules

are very strong. The shape of simple covalent molecules can be predicted by using

different bonding theories.

2.4.1. Lewis Definition of Covalent Bonding An American chemist named Gilbert N. Lewis developed the Lewis bonding theory in

which electrons are represented as dots. The molecules represented are called Lewis

structures or Lewis electron-dot formulas. Today we use Lewis structures to

determine how atoms are arranged in a molecule and to predict the 3D shape of

molecules. Nonmetal atoms form bonds to achieve a Noble Gas electron

configuration. However, instead of taking electrons away from one another to form

ions, they simply share the electrons in a covalent bond. Lewis defined a covalent

bond as an interaction of two atoms which are held together by a pair of electrons

located between the two atoms. The two electrons are shared by the two atoms. The

following generalizations are helpful in constructing Lewis structures:

I. Count the total number of valence electrons present in the compound by adding the

total valence electrons each atom.



II. Determine the number of pairs (single bonds) by dividing the total valence electrons

by two.

III. Predict the central atom based on the electronegative. The least electronegative

atom in a molecule is often located in the central position and the more

electronegative atoms occupy the outer (terminal) positions.

IV. Connect all atoms by drawing single bonds between all atoms, and then distribute

the remaining valence electrons as lone pairs around outer atoms and around the

central atom so each has an octet.

V. If there are not enough electrons for each atom to have an octet, make double or

triple bonds between central atom and surrounding atoms. The great majority of

inorganic compounds can be described by Lewis structures where each atom

achieves a closed shell electronic configuration.

The number of bonds x in a main-group compound other than H2 is predicted as:

Where h is the number of hydrogen atoms, and p is the number of p-block atoms. The

2h+ 8p is the maximum total number of electrons they can have in their valence

orbitals.

Example 1: F2 x = (2 x 8 -14)/2 = 1, single bond

The resulting molecule, written as F-F, demonstrates an important point. It has six

pairs of electrons which are not shared by the atoms. Such pairs of electrons are

called lone pairs.

Example 2: O2 x = (2 x 8 -12)/2 = 2, double bonds

Oxygen needs two electrons to attain the octet structure. This places two shared pairs

between the two oxygen nuclei. Two pairs of electrons between two atoms constitute



two bonds, that is, a double bond. The bond order in the molecule is 2 and O2 is

written as O=O.

Example 3: NF3 x= (4 x 8 -26)/2 = 3 bonds.

Example: CF4

Step1 1C= 4 X 1= 4electrons

4F= 4 x 7= 28es, 28 + 4= 32electrons

Step2 32/2=16 pairs

Step3

C

F

F

F F

Step4 16 pairs -4 pairs= 12 pairs= 24es

C

F

F

F F

The resulting molecule, written as F-F, demonstrates an important point. It has six

pairs of electrons which are not shared by the atoms. Such pairs of electrons are

called lone pairs.

Example: CH2O

Rule 1 Carbon has 4 valence electrons, each hydrogen has 1 valence electron, and

oxygen has 6 valence electrons.Total number of valence electrons : 4 + 2(1) + 6 = 12

Rule 2 Carbon is at the centre of the molecule because it is less electronegative than

oxygen. Hydrogen is always an outer atom and is never at the centre of the molecule.

Rule 3 Add the bonding electrons.This uses 6 of the 12 valence electrons



Rule 4 Add the remaining 6 lectrons to the outer atom. Hydrogen does not need any

more electrons, but Oxygen needs 6 to complete its octet.

Rule 5 There are no valence electrons left to add to the centre

Rule 6 Oxygen shares one of its lone pairs with C and O and give the desired 8

electron total(double bonds).

Exercise: Draw the Lewis structure and find the bonds between the molecules of the

following compounds; ClO4-, SO3, NO2

-, H2O and BF3.

Resonance Structures

We have assumed up to this point that there is one correct Lewis structure. Now we

turn our attention to Resonance structures: the structure the same relative

placement of atoms but different locations of bonding and non-bonding electron pairs.

Systems which have more than one Lewis structure are Resonance Structures.

Resonance is delocalization of electrons through double bond. NO3- is a classic

example of resonance:



Formal Charge: The hypothetical charge on an atom in the molecules. It helps to

check whether the Lewis structure is stable or not. The smallest formal charge is the

more stable in Lewis structure regardless numerical sign. A more negative formal

charge should reside on an atom with a larger EN value.

FC = formal charge; #Ve. = Number of valence electrons

#BE = bonding electrons; #LPE = lone pair electrons

Example: NCO- has three possible resonance forms.

Forms B and C have negative formal charges on N and O. These forms are more

important than Form A. Form C has a negative charge on O which is more

electronegative than N. Therefore, Form C contributes the most to the resonance

hybrid and stable.

Limitations: Lewis theory was good for s-block and p-block elements but not for d-

block elements. Except in simple cases, Lewis structure can predict neither the 3D

shape of the species (bond angles) nor the relative internuclear distances (bond

LPEBE ##

2

1- Ve.# FC



lengths). Lewis theory cannot write one correct structure for many molecules

where resonance is important. Lewis theory often does not predict the correct

magnetic behavior of molecules. Oxygen, O2, is paramagnetic, though the Lewis

structure predicts it is diamagnetic.

Exceptions to the Octet Rule

Electron-Deficient Molecules: gaseous molecules containing either Be or B as the central

atom; have fewer than 8 electrons around the Be or B (4 e- around Be and 6 e- around B) (BF3).

Odd-Electron Molecules: have an odd number of valence electrons; examples include free

radicals, which contain a lone (unpaired) electron and are paramagnetic (use formal charges to

locate the lone electron) (NO2). Example: CH3, OH, H, NO2 etc.

Expanded Valence Shells: for molecules that have more than 8 electrons around the central

atom; use empty outer d orbitals; occurs only with a central atom from Period 3 or higher (SF6,

PCl5). Eg: PCl5, SF6, H2SO4, H3PO4

2.4.2. Valence shell electron Pair repulsion (VSEPR) Model This model provides a method for predicting the shape of molecules, based on the

electron pair electrostatic repulsion. It was described by Sidgwick and Powell in 1940

and further developed by Gillespie and Nyholm in 1957. The VSEPR method predicts

shapes that compare favorably with those determined experimentally. However, this

approach at best provides approximate shapes for molecules, not a complete picture of

bonding. Electrons repel each other because they are negatively charged. The

quantum mechanical rules force some of them to be fairly close to each other in

bonding pairs or lone pairs, but each pair repels all other pairs. The repulsion force

between the lone pairs and bond pairs follows this order: Lone pair - Lone pair > Lone

Pair - Bond pair > Bond pair - Bond pair in a chemical bond.

According to the VSEPR model, therefore, molecules adopt geometries in which their

valence electron pairs position themselves as far from each other as possible. A

molecule can be described by the generic formula AXmEn, where A is the central

atom, X stands for any atom or group of atoms surrounding the central atom, and E

represents a lone pair of electrons. The steric number (SN = m + n) is the number of

positions occupied by atoms or lone pairs around a central atom.

How it works: the following assumptions are important in order to determine the

molecular shape of a compound.



1. Draw the Lewis structure.

2. Modify the Lewis structure by assigning all atoms or groups bonded to the central

atom as single (-), double (=), or triple (≡) bonds.

3. Determine the number of valence electrons in the central atom.

4. Count all σ-bonds around the central atom and the molecular geometry is dictated

by σ-bond electron pairs. When the valence shell of central atom contains only

bond pairs, the molecule assumes symmetrical geometry due to even repulsions

between them. But the symmetry is distorted when there are also lone pairs along

with bond pairs due to uneven repulsion forces.

5. The electrons contributed by the central atom in any π-bonds should be

'discounted' by subtracting one electron for each π-bond from the electron count.

(Considering point 4, one can ignore any double bonds when counting electrons in

VSEPR).

6. If the molecule is charged, the corresponding number of electrons is either added

(for a negative charge) or subtracted (for positive charge) from the electron count.

7. Dividing the resulting number of electrons by 2 gives the number of electron pairs

around the central atom. Now geometry can be assigned.

VESPR predicts the following geometries depending on the lone pairs and bond

pairs as shown below.

VSEPR Examples



Limitations

1. This approach is time-consuming, because it

requires handling all valence electrons.

2. It is easy to make simple arithmetic mistakes when

you work with large numbers.

3. The number of electrons around the central atom in

a molecule can often be ≠ 8 (e.g. 6, 10, 12), that can be

confusing and make you feel unsure.

4. Accurate Lewis structures can be difficult to draw

for many compounds. This can complicate making the

decisions about the shape.

This may be summarized:

1. The shape of the molecule is determined by repulsions between all of the electron

pairs present in the valence shell.

2. A lone pair of electrons takes up more space round the central atom than bond pair

since the lone pair is attracted to one nucleus whilst the bond pair is shared by two



nuclei. It follows that repulsion between two lone pairs is greater than repulsion

between a lone pair and a bond pair, which in turn is greater than the repulsion

between two bond pairs. Thus the presence of lone pairs on the central atom causes

slight distortion of the bond angles from the ideal shape. If the angle between a lone

pair, the central atom and a bond pair is increased, it follows that the actual bond

angles between the atoms must be decreased.

3. The magnitude of repulsions between bonding pairs of electrons depends on the

electronegativity difference between the central atom and the other atoms.

4. Double bonds cause more repulsion than single bonds, and triple bonds cause more

repulsion than a double bond.

Effect of lone pairs: Molecules with four electron pairs in their outer shell are based

on a tetrahedron. In CH4 there are four bonding pairs of electrons in the outer shell of

the C atom, and the structure is a regular tetrahedron with bond angles H-C-H of

109°28°. In NH3 the N atom has four electron pairs in the outer shell, made up of

three bond pairs and one lone pair. Because of the lone pair, the bond angle H-N-H is

reduced from the theoretical tetrahedral angle of 109°28° to 107°48°. In .H2O the atom

has four electron pairs in the outer shell. The shape of the H2O molecule is based on a

tetrahedron with two corners occupied by bond pairs and the other two corners

occupied by lone pairs. The presence of two lone pairs reduces the bond angle further

to 104°27°. CH4: H-C-H = 109º28', NH3: H-N-H = 107º48', H2O: H-O-H = 104º27'

In a similar way, SF6 has six bond pairs in the outer shell and is a regular octahedron

with bond angles of exactly 90°. In BrF5, the Br also has six outer pairs of electrons,

made up of five bond pairs, and one lone pair. The lone pair reduces the bond angles

to 84°30°. Whilst it might be expected that two lone pairs would distort the bond

angles in an octahedron still further, in XeF4 the angles are 90°. This is because the

lone pairs are trans to each other in the octahedron, and hence the atoms have a

regular square planar arrangement.

Effect of electronegativity: NF3 and NH3 both have the same structures based on a

tetrahedron with one corner occupied by a lone pair. The high electronegativity of F

pulls the bonding electrons further away from N than in NH3. Thus, repulsion between

bond pairs is less in NF3 than in NH3. Hence the lone pair in NF3 causes a greater

distortion from tetrahedral and gives a F-N-F bond angle of 102°30°, compared with



107°48° in NH3. The same effect is found in H2O (bond angle 104°27°) and F2O (bond

angle 102°).

2.4.3. Valence Bond theory and Hybridization Brainstorming Questions: what mean by valence bond theory? Describe the main

feature of valence bond theory. List the shortcomings of valence bond theory. What is the

purpose of Hybridization?

The valence bond theory was proposed by Heitler and London to explain the formation

of covalent bond quantitatively using quantum mechanics. Later on, in the 1930s,

Linus Pauling improved this theory by introducing the concept of hybridization.

According to valence bond theory, covalent bonds are formed between the two atoms

by the overlap of half filled valence atomic orbitals of each atom containing one

unpaired electron.

Hybridization: is the intermixing of two or more pure atomic orbitals of an atom with

almost same energy to give same number of identical and degenerate new type of

orbitals. The new orbitals formed are also known as hybrid orbitals. Two types of

covalent bonds are formed based on the pattern of overlapping orbitals. These are:

(i) σ-bond: The covalent bond formed due to overlapping of atomic orbital along the

inter nucleus axis is called σ-bond. It is a stronger bond and cylindrically symmetrical.

Depending on the types of orbitals overlapping, the σ-bond is divided into following

types:



(ii) π-bond: The covalent bond formed by sidewise overlapping of atomic orbitals is

called π- bond. In this bond, the electron density is present above and below the

internuclear axis. It is relatively a weaker bond since the electrons are not strongly

attracted by the nuclei of bonding atoms.

Types of Hybridization

During hybridization, the atomic orbitals with different characteristics are mixed with

each other. Hence there is no meaning of hybridization between same types of orbitals

i.e., mixing of two 's' orbitals or two 'p' orbitals is not called hybridization. However,

orbital of 's' type can mix with the orbitals of 'p' type or of 'd' type. Based on the type

and number of orbitals, the hybridization can be subdivided into following types.

sp Hybridization: Intermixing of one 's' and one 'p' orbitals of almost equal energy to

give two identical and degenerate hybrid orbitals is called 'sp' hybridization. These sp-

hybrid orbitals are arranged linearly at 180o of angle. They possess 50% 's' and 50% 'p'

character.

Example: BeCl2; Be is the central atom in this compound and normally Be has 4

atomic number, then its electronic configuration is 1s2, 2s2. Thus, Be contains 2 e-s at

the ground state before bonding in 2s orbital. But at the exited state (during reaction)



one electron promotes to the empty 2p orbital in order to share these two electrons

with two chlorines.

sp2 Hybridization: Intermixing of one 's' and two 'p' orbitals of almost equal energy to

give three identical and degenerate hybrid orbitals is known as sp2 hybridization. The

three sp2 hybrid orbitals are oriented in trigonal planar symmetry at angles of 120o to

each other. The sp2 hybrid orbitals have 33.3% 's' character and 66.6% 'p' character.

Example: BCl3

Sp3 Hybridization: In sp3 hybridization, one 's' and three 'p' orbitals of almost equal

energy intermix to give four identical and degenerate hybrid orbitals. These four sp3

hybrid orbitals are oriented in tetrahedral symmetry with 109o28' angle with each

other. The sp3 hybrid orbitals have 25% 's' character and 75% 'p' character.



Example: CH4

Based on the type of hybridization we can predict the shape of simple molecules.

Hybridization Shape Example

Sp Linear BeCl2

sp2 Bent BF3

sp3 Tetrahedral CH4

sp3d T. bipyramid PCl5

sp3d2 Octahedral SF6

Limitation of VBT

The theory is unable to adequately explain electronic and magnetic properties of

complexes.

VBT is widely used in organic and main group element chemistry.

In transition metal chemistry VBT is superseded by the Crystal Field Theory (CFT).

2.4.4. Molecular Orbital Theory (MOT) The Molecular Orbital Theory, initially developed by Robert S. Mullikan, incorporates

the wave like characteristics of electrons in describing bonding behavior. In Molecular

Orbital Theory, the bonding between atoms is described as a combination of their

atomic orbitals. While the Valence Bond Theory and Lewis Structures sufficiently

explain simple models, the Molecular Orbital Theory provides answers to more

complex questions. In the Molecular Orbital Theory, the electrons are delocalized since

they are not assigned to a particular atom or bond. Instead, the electrons are

―smeared out‖ across the molecule.

http://chemwiki.ucdavis.edu/Theoretical_Chemistry/Chemical_Bonding/Valence_Bond_Theory



The Molecular Orbital Theory allows one to predict the distribution of electrons in a

molecule which in turn can help to predict molecular properties such as shape,

magnetism, and Bond Order.

Bonding and Antibonding Molecular Orbitals (MOs)

Bonding Molecular Orbital: MOs with electron density concentrated in the regions

between atoms. Bonding MOs have lower energy and greater stability than the atomic

orbitals from which it was formed. Antibonding Molecular Orbital: MOs with electron

density concentrated in regions other than between the atoms. Anitbonding MOs have

higher energy and lower stability than the atomic orbitals from which it was formed.

Linear Combination of Atomic Orbitals (LCAO)

A linear combination of atomic orbitals, or LCAO, is a quantum superposition of

atomic orbitals and a technique for calculating molecular orbitals in quantum

chemistry. This supposes that we can construct molecular orbitals from linear

superposition of atomic orbitals centered on individual atoms. It can be expressed

using the following equation: += ca+a+ cb+b or -= ca-a-cb-b where is the

molecular wave function, a and b are atomic wave functions, and ca, and cb are

adjustable mixing coefficients. The coefficients can be equal or unequal, positive or

negative, depending on the individual orbitals and their energies.

Three conditions are essential for overlap bonding MO.



The symmetry of the orbitals must be such that regions with the same sign of +

overlap.

The energies of the atomic orbitals must be similar. When the energies differ by a

large amount, the change in energy on formation of the molecular orbitals is small

and the net reduction in energy of the electrons is too small for significant bonding.

The distance between the atoms must be short enough to provide good overlap of

the orbitals, but not so short that repulsive forces of other electrons or the nuclei

interfere. When these conditions are met, the overall energy of the electrons in the

occupied molecular orbitals will be lower in energy than the overall energy of the

electrons in the original atomic orbitals, and the resulting molecule has a lower

total energy than the separated atoms.

Types of Bonds in Molecular orbital

Sigma Bonds: Molecular orbitals that are symmetrical about the axis of the bond are

called sigma molecular orbitals, often abbreviated by the Greek letter σ. There are two

types of sigma orbitals formed, antibonding sigma orbitals (abbreviated σ*), and

bonding sigma orbitals (abbreviated σ). In sigma bonding orbitals, the in phase atomic

orbitals overlap end to end causing an increase in electron density along the bond

axis. In sigma antibonding orbitals (σ*), the out of phase overlap of orbitals interfere

destructively which results in a low electron density between the nuclei. In the s-

orbital, the sigma boding orbital can be represented as σs and the antibonding will be

written as σ*s. Sigma bonding orbitals and antibonding orbitals can also be formed

between p orbitals. Sigma molecular orbitals formed by p orbitals are often

differentiated from other types of sigma orbitals by adding the subscript p below it as

σ*2p.

Pi Bonds: The pi bonding bonds as a side to side overlap, which then causes there to

be no electron density along the axis, but there is density above and below the axis. In

the pi bonding, there are an overlap of atomic orbitals in between py-py and pz-pz

orbitals to form π2py, π2pz bonding and π*2py, π*2pz antibonding molecular orbitals. The

2pz-2pz overlap is similar to the 2py-2py overlap because it is just the orbitals of the 2pz

rotated 90 degrees about the axis.

Bond Order: Bond Order indicates the strength of the bond. The higher the Bond

Order the stronger the bond. Bond order can be calculated as follows:



Example: H2, BO =½ (2 –0) = 1 H–H single bond will be formed.

Bond order determines the stability of the molecule. If the Bond Order is Zero,

then no bonds are produced and the molecule is not stable (for example He2). If the

Bond Order is 1, then it is a single covalent bond. The higher the Bond Order, the

more stable the molecule is. An advantage of Molecular Orbital Theory when it comes

to Bond Order is that it can more accurately describe partial bonds (for example in

H2+, where the Bond Order=1/2), than Lewis Structures.

Rules of MOT: MO Theory has five basic rules:

The number of molecular orbitals = the number of atomic orbitals combined.

Of the two MO's, one is a bonding orbital (lower energy) and one is an anti-

bonding orbital (higher energy).

Electrons enter the lowest orbital available.

The maximum # of electrons in an orbital is 2 (Pauli Exclusion Principle).

Electrons spread out before pairing up (Hund's Rule).

Drawing Molecular Orbital Diagrams

Determine the number of electrons in the molecule.

Fill the molecular orbitals from bottom to top until all the electrons are added.

Describe the electrons with arrows. Put two arrows in each molecular orbital, with

the first arrow pointing up and the second pointing down.

Orbitals of equal energy are half filled with parallel spin before they begin to pair

up.

Molecular Orbital Energy level Diagram for Homonuclear and Heteronuclear

Diatomic Molecules

Homonuclear diatomic molecules are molecules made up of the same or identical

elements. E.g., H2, N2, O2, F2, Cl2, Br2, and I2.The molecular electronic configuration of

molecules is determined using spectroscopic data. For period 2 diatomic molecules up

to and including N2 the electronic configuration is: 2s, *2s, π2p, 2p, π*2p, *2p. For

period 2 diatomic molecules O2, F2, and Ne2: π2p and 2p change order and the

electronic configuration is: 2s, *2s, 2p, π2p, π*2p, *2p as shown in the table below.

Hetronuclear diatomic molecules are molecules formed from two different elements. In



heteronuclear diatomic molecules, mixing of atomic orbitals only occurs when the

electronegativity values are similar.

Examples:

Molecule Energy level Diagram

He2, Electron Configuration of He

atom is:

He: 1s2 ,Electrons found in He2 are:

He + He = 2 + 2 = 4

Molecular orbital configuration for

He2: (1s)2(*1s)2,

Bond order for He2: BO = ½ (2 – 2) = 0,

Therefore, He2 does not exist.

He2+, Electron Configuration of He

atom is: 1s2 ,Electrons found in He2+

are:

He+ + He = 1 + 2 = 3

Molecular orbital configuration for

He2+: (1s)2(*1s)1,

Bond order for He2+: BO = ½ (2 –1) = 1,

Therefore the bond between He– He is

single bond.

Lithium (Li2), Electron Configuration of

Li atom is: 1s2 2s1, Electrons found in

Li2 are: Li + Li = 3 + 3 = 6

Molecular orbital configuration for Li2:

(1s)2(*1s)2(2s)2,

Bond order for Li2: BO = ½ (4 – 2) = 1,

Therefore the bond between Li–Li is

single bond.



Nitrogen (N2), Electron Configuration

for N atom is: 1s2, 2s2, 2p3

Electrons found in N2 : N + N = 7 + 7 =

14

MO configuration for N2:

( 1s)2( *1s)2( 2s)2( *2s)2(π2pyz)4( 2px)2

Bond order N2: BO = ½ (10 – 4) = 3

Triple bond is formed b/n N-N atoms

(NN).

Oxygen (O2), Electron Configuration

for O atom is: 1s2, 2s2, 2p4

Electrons found in O2 : N + N = 8 + 8 =

16

MO configuration for O2:

( 1s)2 2( 2s)2( *2s)2( 2px)2(π2py2

π2pz2 ) (π*2py

1 π*2pz1) ( *2px)0

Bond order for O2: BO = ½ (8 – 4) = 2

Double bond is formed b/n O-O atoms

(O=O).

In heteronuclear diatomic molecules, the more electronegative element makes a

greater contribution to the bonding MO’s, and the less electronegative one makes a

greater contribution to the antibonding MO’s. The bonding MO’s have energies

similar to those of the contributing AO’s of the more EN atom. The antibonding MOs

have energies close to those of the contributing AO’s of the less electronegative atom.

NO, Oxygen (NO), Electron

Configuration for: N atom is: 1s2, 2s2,

2p3 and O atom is: 1s2, 2s2, 2p4

Electrons found in NO : N + O = 7 + 8 =

15

MO configuration for NO:

( 1s)2( *1s)2( 2s)2( *2s)2 ( 2px)2(π2py2

π2pz2 ) (π*2py

1 π*2pz0) ( *2px)0

Bond order for NO: BO = ½ (7 – 1) = 3

Triple bond is formed between N-O

atoms (NO).

Exercise: Draw the MO energy level diagram, predict the magnetism, calculate the

band order and determine the stability of the following molecules. NO+, NO-, O2-2, O2

+2,

CO, HF, Be2, C2, H2+ and CO+.



2.5. Structure of Ionic Solids Ionic compounds are formed between highly metallic species and highly

electronegative species (nonmetals) by transfer of electrons from metal to nonmetals

to form solid compounds. Based on the order line of their particles arrangement, we

can divide solids into two broad categories.

Crystalline solids: Solids having well defined shape, their particle-atoms, molecules,

or ions occurs in orderly arrangement are called crystalline solids. “A crystal is a

solid composed of atoms (ions or molecules) arranged in an orderly repetitive array‖.

Example: Diamond. Thus, a crystal may be defined as ―A condition of matter resulting

from an orderly, cohesive, three dimensional arrangements of its component

particles (atoms, ions or molecules) in space‖. This three dimensional arrangement is

called crystal lattice or space lattice. The positions occupied by the particles in the

crystal lattice are called lattice sites or lattice points. The lattices are bound by

surfaces that are usually planar are known as faces of the crystal.

2. Amorphous solids: solids that having poorly defined shapes because they lack

extensive molecular-level ordering of their particles are called amorphous solids/

Non-Crystalline. Amorphous: lacks a systematic atomic arrangement. They have no

long-range order rather random orientations of atoms. Example: SiO2, Glasses

2.5.1. Types of Crystalline Solids

Each type of solids is defined by the type of particles in the crystal, which determines

the force attraction between them.

Atomic solids: individual atoms are held together by dispersion force. Noble gases are

the only example.

Molecular solids: individual molecules are held together by dispersion force, dipole-

dipole or H bonds. Examples, non polar molecules such as Cl2, dry ice (CO2), S8, P4.

Metallic solids: In contrast to the weak dispersion forces between the atoms in atomic

solids, powerful metallic bonding forces hold individual atoms together in metallic

solids. Examples, Na, Zn, Fe

Network Covalent Solids: in such types of solids separate particles are not present.

Instead, strong covalent bonds link the atoms together thought a network covalent

solid. As a consequence of the strong bonding, all these substances have extremely

high melting and boiling points. Examples: SiO2 (quart), C (diamond).



Ionic solids: the positive and negative ions are held together by ion-ion interaction.

The interparticle forces (ionic bonding forces) are much stronger than the van der

Waals forces in atomic or molecular solids. Examples are NaCl, CaF2, and MgO. In

ionic solids to maximize attractions, cations are surrounded by as many anions as

possible, and vice versa, with the smaller of the two ions lying in the spaces (holes)

formed by the packing of the larger. The properties of ionic solids are direct

consequence of the fixed ion positions and very strong interionic forces, which create

lattice energy.

2.5.2. Terminologies

Space lattice: ―A space lattice is an array of points showing how particles (atoms, ions

or molecules) are arranged at different sites in three dimensional spaces.‖ Therefore

space lattice of a crystal has been likened to a wall-paper on which a single pattern is

continuously repeated. Each unit cell requires two vectors ―a‖ and ―b‖ for its

description. A three dimensional space lattice can be similarly divided into unit cells

described by three vectors. The exact location of particles in a unit cell can be

obtained by X-ray diffraction.

The lattice energy of a crystal is the energy evolved when one gram molecule of the

crystal is formed from gaseous ions. The lattice energy is a measure of the strength of

ionic bonds within a specific crystal structure. It is usually defined as the energy

change when a mole of a crystalline solid is formed from its gaseous ions.

M+(g) + X-(g) ∆E MX(s) ∆E = Lattice Energy

Lattice: is any infinite array of point in space, in which each point has identical

surrounding to all other. Regular lattice results in crystal structure portion of regular

lattice at gives unit cell.



Unit cell: The unit cell may be defined as, ―the smallest repeating unit in space

lattice which, when repeated over again, results in a crystal of the given substance‖ .It

is the smallest component of the crystal, which when stacked together with pure

translational repetition reproduces the whole crystal. ―The smallest geometrical

position of the crystal which can be used as repetitive unit to build up the whole

crystal is called a unit cell‖. Generally substances can be packed in four different

types of structures. These are:

simple cubic (SC),

body centered cubic (BCC),

hexagonal closest packed (HCP) and

Cubic closest packed (CCP).

I). Simple Cubic (SC) Packing

It is the simplest structure to visualize since atoms form square packed planes. The

second plane of atoms in this structure stacks directly on top of the first. Each sphere

has the coordination number of six. Simple cubic structures are not the efficient way

of packing atoms, only 52% of the available space is actually occupied by the spheres.





II. Body Centered Cubic (BCC) Packing: is a more efficient way of using space than

simple cubic packing; 68% of the space in this structure is filled. It common

structure for metals in Group IA (Li, K, Na), the heavier metals in Group IIA (Ca, Sr

and Ba) and early transition metals (Ti, V, Cr, Mo and Fe); Each sphere has the

coordination number of eight.



Body-centered cubic unit cells have an atom in the center of the cube as well as one

in each corner. The packing efficiency is 68%, and the coordination number = 8

entered cubic.

NB: both SC and BCC packing structures are called non-close packed structures.



The densest packing is called close packing and consists of layers of atoms at the

apices of a network of hexagons. Close-Packed Structures includes Hexagonal

closest packed (HCP) and cubic closest packed (CCP) are the closest packed

structures. More efficient way of packing spheres, 74% of the space is filled. Small

holes make up the other 26% of the unit cell. Sphere has the coordination number of

in HCP and CCP is 12. In close or closest packing, each metal atom has 12 nearest

neighbors. Six atoms surround an atom in the same plane, and the central atom is

then ―capped‖ by 3 atoms on top and 3 atoms below it. If the bottom ―cap‖ and the

top “cap‖ are directly above each other, in an ABA pattern, the arrangement has a

hexagonal unit cell, or is said to be hexagonal close packed as depicted in figure ―a‖

below. If the bottom and top ―caps‖ are staggered, the unit cell that results is a face-

centered cube as depicted in figure ―b‖ below. This arrangement is called cubic close

packing.

There are two types of holes created by a close-packed arrangement (Octahedral and

tetrahedral holes). Octahedral holes lie within two staggered triangular planes of

atoms. If the larger ions (usually the anions) are in close-packed structures, ions of

the opposite charge occupy these holes, depending primarily on two factors:

1. The relative sizes of the atoms or ions. The radius ratio (usually r+/r- where r+ is

the radius of the cation and r- is the radius of the anion) is generally used to measure

this. Small cations will fit in the tetrahedral or octahedral holes of a close-packed

anion lattice. Somewhat larger cations will fit in the octahedral holes, but not in

tetrahedral holes, of the same lattice. Still larger cations force a change in structure.

2. The relative numbers of cations and anions. For example, a formula of M2X will not

allow a close-packed anion lattice and occupancy of all of the octahedral holes by the

cations because there are too many cations. The structure must have the cations in



tetrahedral holes, have many vacancies in the anion lattice, or have an anion lattice

that is not close-packed.

III. Face-Centered Cubic (FCC) Crystal

Structure

Atoms located at corners and on centers of faces. Many metals, including Ag, Al, Au,

Ca, Co, Cu, Ni, Pb, and Pt, show cubic closest packed structure. Hard spheres touch

along diagonal the cube edge length, a= 2R2. The coordination number, CN =

number of closest neighbors = number of touching atoms, CN = 12. Number of atoms

per unit cell, n = 4. FCC unit cell: 6 face atoms shared by two cells: 6 x 1/2 = 3. 8

corner atoms shared by eight cells: 8 x 1/8 = 1.



IV. Hexagonal close packed: In HCP the spheres in the third plane could pack

directly above the spheres in the first plane to form an ABABAB. .. repeating structure.



HCP important metals such as Be, Co, Mg, and Zn, as well as the rare gas He at low

temperatures show this structure.

2.5.3. Radius Ratio Rule In ideally ionic crystal structures, the coordination numbers of the ions are

determined largely by electrostatic considerations. Cations surround themselves with

as many anions as possible, and vice versa. This maximizes the electrostatic

attractions between neighboring ions of opposite charges and hence maximizes the

lattice energy of the crystal. This requirement led to the formulation of the radius ratio

rule for ionic structure in which the ions and the structure adopted for a particular

compound depend on the relative size of the ions.

Radius ratio rule states:

As the size (ionic radius r) of a cation increases, more anions of a particular size

can pack around it. Thus, knowing the size of the ions, we should be able to

predict a priori which type of crystal packing will be observed. We can account for

the relative size of both ions by using the ratio of the ionic radii: ρ =r+/r−



There are two guidelines to be followed in using this rule. First, a cation must be in

contact with its anionic neighbors. Second, neighboring anions may or may not be in

contact.

Radius ratio used to predict the coordination number and the probable structure.

r+/r- range CN Possible structure Examples

<.155 2 linear BeCl2

0.155-0.225 3 triangular BF3

0.225-.414 4 tetrahedral K2S, ZnS

0.414-0.732 6 octahedral NaCl, TiO2, CdCl2

0.732-0.99 8 Body center CsCl

1 12 FCC/HCP

Examples:

Beryllium sulfide, BeS rBe2+/rS2- = 0.59/1.7 = 0.35 CN = 4

Sodium chloride, NaCl rNa+/rCl- = 1.16/ 1.67 = 0.69 CN= 6

Cesium chloride, CsCl rCl-/ rCs+ = 1.67/1.81 = 0.92 CN = 8

Note: Radius ratio rule gives useful information to predict the coordination number

and structure of ionic compounds, but it does not give the actual structure of the

compound since it is difficult to measure the atomic size accurately due small size.

Example: Zinc sulfide, ZnS r+/r- = 0.52 ∴ C.N. = 6 are predicted…WRONG entire

crystal.





2.5.4. Structures of Binary Compounds

Binary compounds (compounds consisting of two elements) may have very simple

crystal structures and can be described in several different ways.

Ionic compound of the type AX (NaCl, CsCl, ZnS)

Sodium chloride (NaCl): NaCl is made up of face-centered cubes of sodium ions and

face centered cubes of chloride ions combined, but offset by half a unit cell length in

one direction so that the sodium ions are centered in the edges of the chloride lattice

(and vice versa). If all the ions were identical, the NaCl unit cell would be made up of

eight simple cubic unit cells. Many alkali halides and other simple compounds adapt

this structure. But, in these crystals, the ions tend to have quite different sizes,

usually with the anions larger than the cations (each sodium ion is surrounded by six

nearest neighbor chloride ions, and each chloride ion is surrounded by six nearest-

neighbor sodium ions). NaCl is a solid composed of a three-dimensional array of

alternating Na+ ions (green) and Cl− ions (purple) held together by the attraction of

opposite charges.

Crystal structure of sodium chloride (table salt)

Cesium chloride, CsCl



As mentioned previously, a sphere of radius 0.73r will fit exactly in the center of a

cubic structure. Although the fit is not perfect, this is what happens in CsCl, Figure

below where the chloride ions form simple cubes with cesium ions in the centers.

In the same way, the cesium ions form simple cubes with chloride ions in the centers.

The average chloride ion radius is 0.83 times as large as the cesium ion (167 pm and

202 pm, respectively. Only CsCl, CsBr, CsI, TlCl, TlBr and TlI, have this structure at

ordinary temperatures and pressures, although some other alkali halides have this

structure at high pressure and high temperature. The cesium salts can also be made

to crystallize in the NaCl lattice on NaCl or KBr substrates, and CsCl converts to the

NaCl lattice at about 469oC.

Zinc blende, ZnS

In ZnS, the radius ratio of 0.40 suggests a tetrahedral arrangement. Each of Zn2+-ion

is tetrahedrally surrounded by four S2- ions and S2- ion is tetrahedrally surrounded by

four Zn2+ ions. This is a 4:4 arrangement since the coordination number of both ions

is four. Actually two different forms of zinc sulphide exists, 4:4 arrangements, one

being called zinc blended and the other wurtzite.

Zinc blended: is related to a cubic closed-packed structure, while Wurtzite:

hexagonal close-packed structure. Many other ionic compounds, including AgI, CdS,

and Cu (I) halides adopt the zinc blende structure. ZnS has two common crystalline

forms, both with CN 4. Zinc blende is the most common zinc ore and has essentially

the same geometry as diamond, with alternating layers of Zn and S. It can also be

described as having zinc ions and sulfide ions, each in face-centered lattice, so that



each ion is in a tetrahedral hole of the other lattice. The stoichiometry requires half of

these tetrahedral holes to be occupied, with alternating occupied and vacant sites.

zinc blended

wurtzite

Ionic compound of the type AX2

There a two common structure: Fluorite and rutile structures. Many diflourides and

dioxides have one of these structures.

Fluorite (CaF2) structure: Ionic salts with a 1:2; cation: anion ratio, especially those

having relatively large cations and relatively small anions. Example: SrF2 and BaCl2

have the fluorite structure. In fluorite structure each of Ca2+-ion is surrounded by

eight F- ions as Ca2+ ions the coordination number of both ions is not the same, and

four Ca2+ ions are tetrahedral arranged around each F- ion. The coordination number

is eight and four; this arrangement is called an 8:4 arrangement. The fluorite

structure is found when the radius ratio is 0.73 or above. It may relate to the closed-

packed structures.

Antifluorite: The antifluorite structure is often seen in ionic compounds having a

cation: anion ratio of 2:1 and relatively large anion for example, K2S

Fluorite (CaF2) structure b) Rutile structur

Rutile (TiO2) structure



The rutile structure is found where the radius ratio is between 0.73 and 0.41. The

coordination numbers are six and three, each Ti4+-ion being octahedrally surrounded

by six O2- -ions and each O2--ion having three Ti4+-ions round it in a plane triangular

arra022ngement. The rutile structure is not close packed, but the Ti4+-ions may be

considered as forming a considerably distorted body-centered cubic lattice. Rutile

structure: 6:3 coordination; Example: TiO2, GeO2, SnO2, NiF2.

2.6. Structural Defect of Solid Compounds Structural defects are caused by vacancies which is an empty (unoccupied) site of a

crystal lattice, i.e. a missing atom or vacant atomic site such defects may arise either

from imperfect packing during original crystallization or from thermal vibrations of the

atoms at higher temperatures. In the latter case, when the thermal energy due to

vibration is increased, there is always an increased probability that individual atoms

will jump out of their positions. They can move from one site to another more

frequently by displacements of atoms like diffusion, powder sintering, etc. which leave

a holes/vacancies.



In crystal defect, we have two main types of defects (stoichiometric and non-

stoichiometric defects).

Structural Defects

stoichiomeric Defect Non-stoichiometric defect

SchottkyFrenkel

Metal excessMetal defficiency Interstitial

cation & electron

(i) Schottky Defect: a type of vacancies defect that occurred due to the missing of

both positive and negative ions from a crystal. The density of crystal is decreased since

equal numbers of cation and anion are missed by leaving holes. This defect is favored

by high ionic compounds, high coordination number and similar size of cation and

anions. Examples: alkali halides such as NaCl, KI, CsCl and others.

(ii)Frenkel Defect: Frenkel pair/ Frenkel disorder is a type of point defect in

a crystal lattice that formed when smaller ion (usually cation) leaves its place in the

lattice, creating a vacancy (hole), and becomes an interstitial by lodging in a nearby

location. This defect does not have any impact on the density of the solid as it involves

only the migration of the ions within the crystal, thus preserving both the volume as

well as mass. (Frenkel defect is a combination of both vacancy and interstitial defects.

https://en.wikipedia.org/wiki/Crystallographic_defects#Point_defects

https://en.wikipedia.org/wiki/Crystal

https://en.wikipedia.org/wiki/Crystal_structure

https://en.wikipedia.org/wiki/Cation

https://en.wikipedia.org/wiki/Vacancy_defect

https://en.wikipedia.org/wiki/Interstitial_defect



Frenkel defect is arise due to the missing of cations from the lattice site and hereby

occupied the interstitial sites. This defect doesn’t affect the density and volume of

crystals since there is no change particle numbers (i.e. only change the cation position

from crystal line to interstitial position). It favored by low coordination number of

crystal solid and large size difference between cation and anions. Examples: AgCl, AgI

etc.

This type of imperfection is more common highly ionic crystals, because the positive

ions, being smaller in size, get lodged easily in the interstitial positions.

(iii) Interstitial defect: In a closed packed structure of atoms in a crystal if the

atomic packing factor is low, an extra atom may be lodged within the crystal structure.

This is known as interstitial position, i.e. voids. An extra atom can enter the interstitial

space or void between the regularly positioned atoms only when it is substantially

smaller than the parent atoms otherwise it will produce atomic distortion. The defect

caused is known as interstitial defect. In close packed structures, e.g. FCC and HCP,

the largest size of an atom that can fit in the interstitial void or space have a radius



about 22.5% of the radii of parent atoms. Interstitialcies may also be single interstitial,

di-interstitials, and tri-interstitials. We must note that vacancy and interstitialcy are

inverse phenomena.

Interstitial cationInterstitial electron

Interstitial particles

Interstitial defect is occurred when an extra cation and electron are occupied the

interstitial position of lattice structure. This defect is favored by low coordination

number, more covalent character and large size difference between cation and anion

like Frenkel defects.

(iv). Metal deficiency defects: this defect may occur by two ways.

a) Due to missing of cation from the lattice structure which results p-type

semiconductors.

b) By the presence of extra negative ions in the interstitial position of lattice.



(v). Metal excess defect: this defect also arisen by the present of extra cations in

interstitial position or by missing of anion from the lattice structure which left a hole;

this hole is occupied by electron to maintain electrical neutrality as shown below.

2.7. Band Theory of Solids

Bands are distributions of many molecular orbital energy levels, so closely spaced in

energy that they seem to be continuous. Because of the very large number of atoms

that interact in a solid material, the energy levels are so closely spaced that they form

bands. The highest energy filled band, which is analogous to the highest occupied

molecular orbital in a molecule (HOMO), is called the valence band. The next higher

band, which is analogous to the lowest unoccupied molecular orbital (LUMO) in a

molecule, is called the conduction band. The energy separation between these bands

is called the energy gap/band energy.



In conductors, there is no band energy since there is overlapping between valence and

conduction bands. Thus, electrons are freely moved between valence and conduction

bands which results conductivity of the solids usually the metals. In semiconductors,

the band energy is some extent small at where a few electrons can jump between

valence and conduction bands. These result partial conductivity of semiconductors.

But in insulator like plastics, the gab energy between valence and conduction band is

so large; so electrons cannot reach to the conduction band.

The filling of these bands and the size of the energy gap determine if a material is a

conductor (a metal), a semiconductor, or an insulator. In metals there is no energy gap

between filled and unfilled energy levels. A significant number of electrons are

thermally excited into empty levels, creating holes in the filled band. The electrons in a

conduction band and the holes in a valence band can move throughout the material,

allowing it to easily conduct electricity. In semiconductors gab energy is small, but

large enough so that a fairly small number of electrons are in the conduction band

due to thermal energy, and these materials conduct poorly. In insulators gab energy is

large so that electrons are not promoted to the conduction band due to thermal

energy, and these materials do not conduct electricity. The level of the highest

occupied molecular orbital at absolute zero refers Fermi level. It is usually found at

the center between the valence and conduction bands. The particles in this state each

have their own quantum states and generally do not interact with each other. When



the temperature begins to rise above absolute zero, these particles will begin to occupy

states above the Fermi level and states below the Fermi level become unoccupied.

Exercise: Explain how to enhance the conductivity of semiconductor?

Unit Summary Atoms are combined in order to have lower energy, attain octet rule and have net

attraction force. These proceed by losing, gaining or by sharing their valence electrons.

Based on these interaction chemical bonds might be ionic (metal and nonmetal

interaction, metallic bond (metal-metal interaction or covalent (nonmetal- nonmetal

interaction). Beyond these hydrogen and dative bond are a special type of covalent

bond since the former one is formed between hydrogen and small size and high

electronegative nonmetals(F,O.N) without sharing of electrons whereas the later one is

formed a lone pairs of electrons are originated from a single Lewis base and final

aggregate to form and adduct. Scholars have developed different bonding theories to

explain the nature of bond and stability. These are Lewis bonding theory (used for

molecules that satisfy octet rules; valence bond theory (half filled –half filled interaction

of AOs); hybridization (mixing of different AOs with equivalent energy levels which

proceed through ground state, excitation and mixing steps); VSEPRT which deals

arrangement of molecules around central elements or arrangements of both bonding

and lone pair electrons, we should put them far apart to minimize the repulsion forces

among them. A molecular orbital describes a region of space in a molecule where

electrons are most likely to be found, and it has a specific size, shape, and energy level

which formed by linear combination of atomic orbitals (LCAO) on different atoms( the

No. of MOs that formed always equal to the no. of AOs available). This indicates when

two AOs interact; two MOs are formed i.e. bonding MOs that are lower in energy and

anti-bonding MOs that are higher in energy. Based on MOT we can determine the bond

order, stability, bond existence and magnetic nature of molecules. This chapter also

discussed about structure of ionic solids, close packing, radius ration and classification

of ionic solids in the form of AX and AX2. In last section we have seen about structural

defects of solid compounds such as Sckotty defect (missing of both cation and anions

from the structure due to equivalent size of cation and anion); Frenkel defect (occurs

due to small size of cation missing from the structural arrangement); intestinal defect

(extra cation and extra electron are occupied the interstitial position of lattice structure);

metal deficiency ( occurs either missing of cation or due to the presence extra negative



ions); metal excess (extra cations in interstitial position or by missing of anion from the

lattice structure which left a hole; this hole is occupied by electron). Solid materials are

classified into conductor (have equivalent balance band and conduction band);

metalloid (have some gab energy between balance band and conduction band) or

insulator (have high gab energy between balance band and conduction band).

Review Questions 1. Do electron dot structures accurately describe the shapes of molecules? Explain your

answer.

2. Draw Lewis structure for N2O by showing all possible Lewis steps and determine the

most stable one from all possible resonance structures?

3. Consider the bonding in the molecule O2-2, O2, O2

+2 and C2+2

A. Determine the best a Lewis structure

B. Sketch an MO energy level diagram.

C. Write valence MO electronic configuration.

D. Number of electrons on bonding MO and anti-bonding?

E. Determine HOMO and LUMO’s

F. Bond orders of each molecules

G. Which oxygen molecule has highest bond order?

H. Which molecule/s is/are paramagnetic which are not?

4. What crystal structure would you predict for a crystal containing Fe2+ and F−? The radii

are 75 and 136 pm, respectively

5. Define the following terms.

A. Unit cell

B. Lattice point

C. Crystal structure

D. Radius ratio

6. What mean by ionic solids? And give it good examples.

7. Mention and explain examples of ionic compounds of the form AX and AX2

8. Determine the hybridization and structure of the following molecules.

A. SO2 B. H2O C. NH3 D. CH4

9. Mention the types of chemical bonding with their specific examples.

10. Show the structure of the following compounds according to Lewis dot structure.

A. SF6 B. XeF4 C. BF3 D. NO2



11. Calculate the bond order and molecular electronic configuration for(O2, O2-,O2

-2).

3. Acid-Base Theory

At the end of this chapter students, will be able to:

Describe the nature of acids and bases.

Describe the overview of acid -base theory.

Explain HSAB principle.

Strength of binary and oxoacids


What is an acid in your definition? What are a base and an acid in Arrhenius,

Bronsted- Lowry, Lewis and Pearson definitions’?

3.1. Introduction Properties of Acids and Bases: Some of the properties of acids and bases are

tabulated in the table shown below.

Property ACIDS BASES

Taste Sour bitter

Color change Blue to red Red to blue

pH <7 >7

Reactivity Reacts with bases to

produce salt and water

Reacts with acids to produce

salt and water

Reactivity with water Produce H+ or H3O+ ions Produce OH- ion

3.2. Terminologies

A. Strong Acids: acids that ionized (dissociated) completely in water. Example:

HClO4, HI, ClO3, HNO3, HBr, H2SO4

B. Strong base: A base that is completely dissociated in water (highly soluble).

Example: Group 1A metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) and

Heavy Group 2A metal hydroxides [Ca(OH)2, Sr(OH)2, and Ba(OH)2]

C. Neutralization Reactions: The neutralization reaction between strong acids and

strong bases has the net ionic reaction. H3O+ (aq)+ OH– (aq) → 2H2O

D. Bystander ions/ spectator ions: They do not participate /appear in the over al

reaction.

E. Protic solvents: solvate anions (negatively charged solutes) strongly via hydrogen

bonding. Water is a protic solvent. Aprotic solvents such as acetone or

dichloromethane tend to have large dipole moments (separation of partial positive

and partial negative charges within the same molecule) and solvate positively

charged species via their negative dipole.

http://en.wikipedia.org/wiki/Anion

http://en.wikipedia.org/wiki/Hydrogen_bonding

http://en.wikipedia.org/wiki/Hydrogen_bonding

http://en.wikipedia.org/wiki/Acetone

http://en.wikipedia.org/wiki/Dichloromethane

http://en.wikipedia.org/wiki/Molecular_dipole_moment



F. Amopteric: chemical species that acts either an acid or a basse. Example: HSO4-,

H2O, H2PO4-, HPO4

2-

G. Autodissociation: any solvent that can dissociate into a cation and an anion.

3.3. Acid-Base Theory

Many acid-base definitions have been proposed and have been useful in particular

types of reactions. Among these are the ones attributed to Arrhenius (based on

hydrogen and hydroxide ion formation), Bronsted-Lowry (hydrogen ion donors and

acceptors), and Lewis (electron pair donors and acceptors) and others several theories.

In chemistry, acids and bases have been defined differently by different sets of

theories. The main concepts or theories that give the definition of acids and bases are

discussed below.

3.3.1. Arrhenius Theory Acid-base chemistry was first satisfactorily explained in molecular terms after Ostwald

and Arrhenius (a Swedish chemist, established the existence of ions in aqueous

solution in 1880-1890 (after much controversy and professional difficulties, Arrhenius

received the 1903 Nobel Prize in Chemistry for this theory). As defined at that time,

Arrhenius acids are substances that form hydrogen ions (H+) or hydronium ions, H3O+

in aqueous solution. An acid is a substance that increases the H+ (or H3O+)

concentration in an aqueous solution. HCl + H 2O H3O+ + Cl- Arrhenius bases are

substances that form hydroxide ions (OH-) in aqueous solution. A base is a substance

that increases the OH- concentration in an aqueous solution. Example: NaOH (aq)

→Na+ (aq) +OH− (aq). The reaction between hydrogen ions and hydroxide ions to form

water is the universal aqueous acid-base reaction. The ions accompanying the

hydrogen and hydroxide ions form a salt, so the overall Arrhenius acid-base reaction

can be written as: Acid + base → salt + water

For example: Hydrochloric acid + Sodium hydroxide →Sodium chloride + Water; H+ +

Cl- + Na+ + OH-→NaCl + H2O

This explanation works well in aqueous solution, but it is inadequate for non-

aqueous solutions and for gas and solid phase reactions in which H+ and OH- may not

exist. The theory defined an acid or a base in terms of aqueous solution, not include

non-aqueous solvents. What about Na2CO3????



3.3.2. Bronsted-Lowry Theory In 1923, Bronsted and Lowry defined an acid as a species with a tendency to lose a

hydrogen ion and a base as a species with a tendency to gain a hydrogen ion. This

definition expanded the Arrhenius list of acids and bases to include the gases HC1

and NH3, along with many other compounds. This definition also introduced the

concept of conjugate acids and bases, differing only in the presence or absence of a

proton, and described all reactions as occurring between a stronger acid and base to

form a weaker acid and base:

H3O+ + NO2- → H2O + HNO2

Acid1 base2 base1 acid2

In water, HCl and NaOH react as the acid H3O+ and the base OH- to form water, which

are the conjugate base of H3O+ and the conjugate acid of OH-. Reactions in non-

aqueous solvents having ionizable hydrogen parallel those in water. An example of

such a solvent is liquid ammonia, in which NH4Cl and NaNH2 react as the acid NH4+

and the base NH2- to form NH3, which are both a conjugate base and a conjugate acid:

NH4+ + Cl- + Na+ + NH2

- → Na+ + Cl- + 2NH3

With the net reaction NH4+ + NH2

- → 2NH3

Acid base conjugate acid and conjugate base

In any solvent, the direction of the reaction always favors the formation of weaker

acids or bases than the reactants. In the two examples above, H3O+ is a stronger acid

than HNO2 and the amide ion is a stronger base than ammonia and ammonium ion is

a stronger acid than ammonia, so the reactions favor formation of HNO2 and ammonia.

In contrast to the Arrhenius definition, the Brønsted-Lowry definition refers to the

products of an acid-base reaction as conjugate acids and bases to refer to the relation

of one proton, and to indicate that there has been a reaction between the two

quantities, rather than a "formation" of salt and water, as explained in the Arrhenius

definition:

AH + B → BH+ + A-

BaseAcidAcidBase

NH4+ + OH-

NH3 + H2O



―Conjugate‖ acids and bases are found on the products side of the equation. A

conjugate base is the same as the starting acid minus H+.

In general acid, base, conjugates can be donated as

Excise: Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-

base pairs:

HC2H3O2(aq) + H2O(l) C2H3O2–(aq) + H3O+(aq)

OH –(aq) + HCO3–(aq) CO3

2–(aq) + H2O(l)

This theory works for all protic solvents (acetic acid, water, liquid ammonia, etc.), not

only for water as Arrhinius’ theory. But it does not explain the acid-base behavior in

aprotic solvents such as benzene and dioxane. Bronsted-Lowery’s theory cannot

explain the reactions between acidic oxides. For instance, CO2, SO2, SO3, etc… and

also the basic oxides like CaO, BaO, MgO, etc. which also take place even in the

absence of the solvent as shown in this reaction. CaO +SO3 CaSO4 (There is no

proton transfer in this example). The theory works very nicely in all protic solvent, but

fails to explain acid-base behavior in aprotic solvents. Finally, this theory keep silent

from giving any idea for substances like BF3, AlCl3 etc…that do not have hydrogen

and cannot donate proton but are known to act as acid. The

+

Cl H

H

H

O

+

H

H

H O

Cl



3.3.3. Lewis Theory Lewis defined a base as an electron-pair donor and an acid as an electron-pair

acceptor. This definition further expands the list to include metal ions and other

electron pair acceptors as acids and provides a handy framework for non-aqueous

reactions. The Lewis definition of acid base reactions, devised by Gilbert N. Lewis in

1923 is an encompassing theory to the Brønsted-Lowry and solvent-system definitions

with regards to the premise of a donation mechanism, which conversely attributes the

donation of electron pairs from bases and the acceptance by acids, rather than

protons or other bonded substances and spans both aqueous and non-aqueous

reactions

Lewis acid - a substance that accepts an electron pair; Lewis base - a substance that

donates an electron pair

In addition to all the reactions discussed previously, the Lewis definition includes

reactions such as:

with silver ion (or other cation) as an acid and ammonia (or other electron-pair donor)

as a base. In reactions such as this one, the product is often called an adduct, a

product of the reaction of a Lewis acid and base to form a new combination. Another

example is the boron trifluoride-ammonia adducts, BF3.NH3. The lone pair in the

HOMO of the ammonia molecule combines with the empty LUMO of the BF3, which

has very large, empty orbital lobes on boron, to form an adduct. Lewis acid-base

adducts involving metal ions are called coordination compounds (bonds formed with

both electrons from one atom are called coordinate bonds).

3.3.4. Lux-Flood theory This theory proposed by German chemist Hermann Lux in 1939, further improved by

Håkon Flood circa 1947 and is still used in modern geochemistry and

electrochemistry of molten salts. This definition describes an acid as an oxide ion

(O2−) acceptor and a base as an oxide ion donor. For example:

MgO (base) + CO2 (acid) → MgCO3

CaO (base) + SiO2 (acid) → CaSiO3

HH

H

BASE

••••••

O—HO—H

H+

ACID

http://en.wikipedia.org/wiki/Gilbert_N._Lewis

http://en.wikipedia.org/wiki/1923

http://en.wikipedia.org/wiki/Hermann_Lux

http://en.wikipedia.org/wiki/H%C3%A5kon_Flood

http://en.wikipedia.org/wiki/Geochemistry

http://en.wikipedia.org/wiki/Electrochemistry

http://en.wikipedia.org/wiki/Molten_salt



NO−3 (base) + S2O2−

7 (acid) → NO+2 + 2 SO2−

4

3.3.5. Usanovich definition Usanovich's theory can be defined as an acid is anything that accepts negative species

(electrons) or donates positive ones, and a base as the reverse. This theory limited only in

redox (oxidation-reduction) reactions as a special case of acid-base reactions. Some

examples of Usanovich acid-base reactions include:

Na2O (base) + SO3 (acid) → 2 Na++ SO2−4 (species exchanged: anion O2−)

3 (NH4)2S (base) + Sb2S3 (acid) → 6 NH+4 + 2SbS3−

4 (species exchanged: anion S2−)

Na (base) + Cl (acid) → Na+ + Cl−(species exchanged: electron)

3.3.6. Solvent System Definition of Acids and Bases The solvent system definition applies to any solvent that can dissociate into a cation

and an anion (auto-dissociation), where the cation resulting from auto-dissociation of

the solvent is the acid and the anion is the base. Solutes that increase the

concentration of the cation of the solvent are considered as acids whereas solutes

that increase the concentration of the anion are considered as bases. A solute causing

an increase in the concentration of the solvonium ions and a decrease in the solvate

ions is an acid and one causing the reverse is a base. For any solvent that can

dissociate into a cation and an anion, the cation is the acid, and the anion is the base.

The classic solvent system is water, which undergoes auto-dissociation:

By the solvent system definition, the cation, H3O+ is the acid and the anion, OH-, is

the base. For example, in the reaction

Sulfuric acid increases the concentration of the hydronium ion so it is an acid. The

solvent system approach can also be used with solvents that do not contain hydrogen.

For example, BrF3 also undergoes auto-dissociation:

Solutes that increase the concentration of the acid, BrF2+, are considered acids. For

example, SbF5 is an acid in BrF3:

SbF5 + BrF3→ BrF2+ + SbF6

-

and solutes such as KF that increase the concentration of BrF4- are considered bases:

F- + BrF3 →BrF4-

http://en.wikipedia.org/wiki/Redox



The Arrhenius, Bronsted-Lowry, and solvent system neutralization reactions can be

compared as follows:

Arrhenius: acid + base →salt + water

Bronsted: acid 1 + base 2 →base 1 + acid 2

Solvent System: acid + base →Solvent

This definition is based on a generalization of the earlier Arrhenius definition to all

autodissociating solvents. In all such solvents there is a certain concentration of a

positive species, solvonium cations and negative species, solvate anions, in

equilibrium with the neutral solvent molecules. For example:

2H2O ⇌ H3O+ (hydronium) + OH- (hydroxide)

2NH3 ⇌ NH4+ (ammonium) + NH2

− (amide) or even some aprotic systems:

N2O4 ⇌ NO+ (nitrosonium) + NO3− (nitrate)

For example, water and ammonia undergo such dissociation into hydronium and

hydroxide, and ammonium and amide, respectively:

2H2O H3O++ OH− 2 NH3 NH+ 4 + NH−

3.3.7. Hard - Soft Acids – Bases/Pearson's (HSAB) Principle Ralph Pearson developed the Type A and and Type B acid-base classification by

explaining the differential complexation behaviour of cations and ligands in terms of

electron pair donating Lewis bases and electron pair accepting Lewis acids: Lewis

acid + Lewis base Lewis acid/base complex

Type A metal cations included: Alkali metal cations: Li+ to Cs+; Alkaline earth metal

cations: Be2+ to Ba2+; Lighter transition metal cations in higher oxidation states: Ti4+,

Cr3+, Fe3+, Co3+ , and the proton, H+

Type B metal cations include: Heavier transition metal cations in lower oxidation

states: Cu+, Ag+, Cd2+, Hg+, Ni2+, Pd2+, Pt2+.

In 1963, Ralph G. Pearson, an inorganic chemistry professor at Northwestern

University, first used the adjectives hard soft to describe sets of Lewis acids and bases

that had been segregated according to their characteristics. Pearson can classify acids

and bases into hard soft acid base (HSAB) based on:

The size of cation and anion

The charge of cation and anion

The polarizability of anion and the polarizing power of cation

http://en.wikipedia.org/wiki/Hydronium

http://en.wikipedia.org/wiki/Hydroxide

http://en.wikipedia.org/wiki/Ammonium

http://en.wikipedia.org/wiki/Amide

http://en.wikipedia.org/wiki/Nitrosonium

http://en.wikipedia.org/wiki/Nitrate

http://en.wikipedia.org/wiki/Ammonia

http://en.wikipedia.org/wiki/Hydronium

http://en.wikipedia.org/wiki/Hydroxide

http://en.wikipedia.org/wiki/Ammonium

http://en.wikipedia.org/wiki/Metal_amides#Alkali_metal_amides

http://www.chem.ucsb.edu/people/faculty/pearson/index.shtml



A. Hard acids are cations that have small size and high positive charged species which

result less polarizable by anions. The hard acids are therefore any cations with large

positive charge (3+ or larger) or those whose d electrons are relatively unavailable for

π bonding (e.g., alkaline earth ions, Al3+). Other hard acid cations that do not fit this

description are Cr3+, Mn2+, Fe3+ and, Co3+.

B. Soft acids are anions that have large size and low positively charged species (they

have low polarizing power). Polarizing power = the capability of cations to distort

other charged ions (mainly aions). Example: most transition metals with low charge

density, such as Ag+, is considered as soft acids. Soft acids are those whose d

electrons or orbital’s are readily available for π bonding (+1 cations, heavier +2

cations). In addition, the larger and more massive the atom, the softer it is likely to

be, because the large numbers of inner electrons shield the outer ones and make the

atom more polarizable

C. Hard bases anions with small size and low negative ions. They are less polarizable

species). Examples: Cl-, F-, O2- etc. Polarizable = easily distorted by other charged

ions (mainly by cations).

D. Soft base are anions with large size and high negative ions. They are highly

polarizable species.

Pearson (HSAB) principle used to predict:

the stability of a given product

the reaction direction

Solubility of the compound

Hard acids prefer to associate/coordinate with hard bases, and soft acids prefer soft

bases. The polarizability of an acid or base plays a role in its reactivity. Hard acids

and bases are small, compact, and non-polarizable whereas Soft acids and bases are

larger, with a more diffuse distribution of electrons. Hard/Hard and Soft/Soft

interactions are the most favorable= Hard acids react preferentially with hard

bases, and soft acids react preferentially with soft bases.

Example: Ag+ is large and polarizable = Soft; Softness of Halides: I- > Br- > Cl- > F-.

AgI has the strongest interaction, thus the lowest solubility. Softness is also

associated with covalent bonds, not ionic bonds.

Hard acids or bases are compact, with the electrons held fairly tightly by the



nucleus. They are not very polarizable. F- is a hard base, and metal ions such as Li+,

a hard acid. Large, highly polarizable ions are categorized as ―soft.‖ Iodide is a soft

base; and. The halide ions range from F-, a very hard base, through less hard C1- and

Br- to I-, a soft base. Reactions are more favorable for hard-hard and soft-soft

interactions than for a mix of hard and soft in the reactants. For example, in

aqueous solution:

Ag+ + I- →AgI(s) is a very favorable soft-soft reaction; AgI is very insoluble.

Li+ + F- →LiF(s) is a very favorable hard-hard reaction; LiF is only slightly soluble.

Ag+ + F- →AgF(s) is a soft-hard reaction that is not favored; AgF is moderately soluble.

Li+ + I- →LiI(s) is a hard-soft reaction that is not favored; LiI is soluble.

Bear in your mind!! There is no pure hard soft acid base classification but we can

classify in relative way based on their size, charge and polarization.

Table 3.1: Hard and Soft bases

Table 3.2: Hard and Soft Acids

Example, S2- is softer than O2- because it has more electrons spread over a slightly

larger volume, making S2- more polarizable. Within a group, such comparisons are

easy; as the electronic structure and size change; comparisons become more difficult

but are still possible. Thus, S2- is softer than C1-, which has the same electronic

structure, because S2- has a smaller nuclear charge and a slightly larger size. As a



result, the negative charge is more available for polarization. Soft acids tend to react

with soft bases and hard acids with hard bases, so the reactions produce hard-hard

and soft-soft combinations. Example: Is OH- or S2- more likely to form insoluble salts

with 3+ transition metal ions? Which is more likely to form insoluble salts with 2+

transition metal ions?

Because OH- is hard and S2- is soft, OH- is more likely to form insoluble salts with 3+

transition metal ions (hard) and S2- is more likely to form insoluble salts with 2+

transition metal ions (borderline or soft).

More detailed comparisons are possible, but another factor, called the inherent acid-

base strength, must also be kept in mind in these comparisons. An acid or a base may

be either hard or soft and at the same time be either strong or weak. The strength of

the acid or base may be more important than the hard-soft characteristics; both must

be considered at the same time. If two soft bases are in competition for the same acid,

the one with more inherent base strength may be favored unless there is considerable

difference in softness. As an example, consider the following reaction. Two hard-soft

combinations react to give a hard-hard and a soft-soft combination, although ZnO is

composed of the strongest acid (Zn2+) and the strongest base (O2-):

In this case, the HSAB parameters are more important than acid-base strength,

because Zn2+ is considerably softer than Li+. As a general rule, hard-hard

combinations are more favorable energetically than soft-soft combinations.

Soft acids are defined as electron-pair acceptors--Lewis acids--in which the acceptor

atom has a zero or low positive charge and a relatively large size. These characteristics

give rise to low electronegativity and high polarization of valence electrons, meaning

soft acids are easily oxidized. Hard acids have the opposite characteristics that result

in low polarization of valence electrons.

3.4. Strength of Binary Acids Binary acids are certain molecular compounds in which hydrogen is combined with a

second nonmetallic element. Examples: HF, HCl, HBr, HI. The names of binary acids

begin with hydro- followed by the name of the other element modified to end with -ic.

The strength of the acid is measured by its tendency to ionize: HA H+ + A- . Two

http://en.wikipedia.org/wiki/Hydrogen

http://en.wikipedia.org/wiki/Nonmetal

http://en.wikipedia.org/wiki/Chemical_element



factors influence the extent to which the acid undergoes ionization. One is the

strength of the H-A bond, the stronger the bond, the more difficult it is for the HA

molecule to break up and hence the weaker the acid. The other factor is the polarity

of the H-A bond. If the bond is highly polarized, that is if there is a large accumulation

of positive and negative charges on the H and A atoms, HA will tend to break up into

H+ and A- ions. So a high degree of polarity characterizes a stronger acid. The

halogens form a series of binary acids called the hydrohalic acids. The strengths of

the hydrohalic acids increase in the following order: HF << HCl < HBr < HI. HF

requires highest bond dissociation energy (568.2 kJ) to break the H-F bond, HF is a

weak acid. At the other extreme in the series, HI has the lowest bond energy (298.3

kJ), so HI is the strongest acid of the group. In this series of acids the polarity of the

bond actually decreases from HF to HI. This property should enhance the acidity of

HF relative to the other acids in the series, but its magnitude is not great enough to

overcome the trend in bond dissociation energies.

In any vertical column (Group) of nonmetallic elements, there is a tendency toward

increasing acidity of the hydride with increasing atomic number (as you go down the

group). For example, among the group VIA elements the acid strength increases in

the order: H2O< H2S<H2Se<H2Te. This order arises primarily because the bond

energies steadily decrease in this series as the central atom grows larger and the

overlaps of atomic orbitals grow smaller, just as in the case of the hydrides of the

halogens above.

3.5. Strength of Oxyacids

Oxyacids: any oxygen-containing acid which expresses them by the general formula

HaXbOc, with X representing an element other than hydrogen or oxygen. The strength

of an oxyacid is defined by the extent to which it dissociates in water (i.e., its ability to

form H+ ions). In general, the relative strength of oxyacids can be predicted on the

basis of the electronegativity and oxidation number of the central nonmetal atom (the

number of O atoms).

The acid strength increases as the electronegativity of the central atom increases.

For example, because the electronegativity of chlorine (Cl) is greater than that of sulfur

(S), which is in turn greater than that of phosphorus (P), it can be predicted that

perchloric acid, HClO4, is a stronger acid than sulfuric acid, H2SO4, which should be a

stronger acid than phosphoric acid, H3PO4. For a given nonmetal central atom, the



acid strength increases as the oxidation number of the central atom increases. For

example, nitric acid, HNO3, in which the nitrogen (N) atom has an oxidation number of

+5, is a stronger acid than nitrous acid, HNO2, where the nitrogen oxidation state is

+3. In the same manner, sulfuric acid, H2SO4, with sulfur in its +6 oxidation state, is

a stronger acid than sulfurous acid, H2SO3, where a +4 oxidation number of sulfur

exists.

We can summarize these ideas as two simple rules that relate the acid strength of

oxyacids to the electronegativity of X and to the number of groups attached to X:

1. For oxyacids that have the same number of oxygen atoms, acid strength increases

with increasing electro-negativity of the central atom X. This happens for two

reasons: First, the O-H becomes more polar, thereby favoring loss of H+. Second,

because the conjugate base is usually an anion, its stability generally increases as the

electro-negativity of X increases.

2. For oxyacids that have the same central atom X, acid strength increases as the

number of oxygen atoms attached to X increases.

Many oxyacids, H2SO4 for example, contain additional oxygen atoms bonded to the

central atom X. The additional electronegative oxygen atoms pull electron density

from the O-H bond, further increasing its polarity. Increasing the number of oxygen

atoms also helps stabilize the conjugate base by increasing its ability to spread out its

negative charge. The salt of an oxyacid is a compound formed when the acid reacts

with a base: acid + base → salt + water. This type of reaction is called neutralization,

because the solution made is neutral.

In the series of oxyacids of chlorine, the acid strength in aqueous solution is in the

order: HClO4 > HClO3 > HClO2 > HClO

Pauling suggested a rule that predicts the strength of such acids semi-quantitatively,

based on n, the number of non-hydrogenated oxygen atoms per molecule. The

molecular explanation for these approximations hinges on electro negativity. Each

non-hydrogenated oxygen atom is highly electronegative; it draws electrons away from

the central atom, increasing the positive charge on the central atom. This positive

charge in turn draws the electrons of the hydrogenated oxygen toward itself. The net



result is a weaker O - H bond (lower electron density in these bonds), which makes it

easier for the molecule to act as an acid by losing the H+. As the number of highly

electronegative oxygens increases, the acid strength of the molecule also increases.

The same argument can be seen from the point of view of the conjugate base. The

negative charge of the conjugate base is spread over all the non-hydrogenated

oxygen’s. The larger the number of these oxygen’s to share the negative charge, the

more stable and weaker the conjugate base and the stronger the hydrogenated acid.

This explanation gives the same result as the first: the larger the number of non-

hydrogenated oxygen atoms, the stronger the acid.

3.6. pH scale and pH calculation The pH scale is a way of expressing the strength of acids and bases. Instead of using

very small numbers, we just use the negative power of 10 on the Molarity of the H+ (or

OH-) ion. Under7=acid, 7=neutral, Over 7 = base

pH = - log [H+] (Remember that the [ ] mean Molarity)

A solution has a pH of 8.5. What is the Molarity of hydrogen ions in the solution?

• pH = - log [H+]

• 8.5 = - log [H+]

• -8.5 = log [H+]

• Antilog -8.5 = antilog (log [H+])

• 10-8.5 = [H+]

• 3.16 X 10-9 = [H+]

• Kw = [H3O+] [OH-] = 1.00 x 10-14 at 25 oC

• In a neutral solution [H3O+] = [OH-]

• so Kw = [H3O+]2 = [OH-]2

• Since acids and bases are opposites, pH and pOH are opposites!

• pOH does not really exist, but it is useful for changing bases to pH.

• pOH looks at the perspective of a base

• pOH = - log [OH-]

• Since pH and pOH are on opposite ends,

• pH + pOH = 14

• d so [H3O+] = [OH-] = 1.00 x 10-7 M

What is the pH of the 0.0010 M NaOH solution?

[OH-] = 0.0010 (or 1.0 X 10-3 M)



pOH = - log 0.0010

pOH = 3

pH = 14 – 3 = 11

OR Kw = [H3O+] [OH-]

[H3O+] = 1.0 x 10-11 M

pH = - log (1.0 x 10-11) = 11.00

The pH of rainwater collected in a certain region of the northeastern United States on

a particular day was 4.82. What is the H+ ion concentration of the rainwater? A

chemist dilutes concentrated hydrochloric acid to make two solutions: (a) 3.0 M and

(b) 0.0024 M. Calculate the [H3O+], pH, [OH-], and pOH of the two solutions at 25°C.

What is the [H3O+], [OH-], and pOH of a solution with pH = 3.67? Is this an acid, base,

or neutral?

Unit Summary Different chemists give their definition for acid- base theories in different time

sequences. Arrhenius concept: Acids are substances which produce hydrogen ions in

aqueous solution while Bases are substances which produce hydroxide ions in

aqueous solution. The theory defined an acid or a base in terms of aqueous solution,

not include non-aqueous solvents. According to Bronsted-Lowry theory, an acid is a

proton donor to any other substance and a base as a proton acceptor a proton from

any other substance. During gaining or losing of a proton, a conjugate acid-base pairs

are formed; strong acid form weak conjugate base and vice versa. According to Lewis

concept, a base is defined as a substance which can donate a pair of electrons

whereas an acid is a substance which can accept a pair of electrons. Lewis

theory is wider and inclusive. Pearson classifies Lewis acid/ base into soft/ hard acid-

base based on the size of cation and anion, charge of caion and anions. Soft

bases( Lewis bases with large size and high negative ion that leads high polarizability);

they are donor atoms relatively low electro negativity and high polarizability and

are easy to oxidize. High polarizability: they hold their valence electrons loosely. Hard

bases (Lewis bases with small size and low negative ion that leads less polarizability);

they donor atoms which have high electro negativity and low polarizability and are

hard to oxidize. They hold their valence electrons tightly. Soft acids (the acceptor

atoms are large, have low positive charge, and contain unshared pair of

electrons (p or d) in their valence shells). They have low polarizing power and



small size. Hard acids (the acceptor atoms with small size, high positive charge,

and high polarizing power/high electro negativity. Soft acids are Lewis acids which

are comparatively larger in size and whose electron clouds are easily polarizable.

These are mostly heavy metal ions generally associated with low (or even zero) positive

oxidation state. The strength of acid base can be determined by different

measurements such as degree of dissociation (strong acid-base highly dissociable);

pH/pOH value (strong acid has low pH value vice versa); concentration of H+/OH- ion

(strong acid/base has high concentration of H+/OH-, respectively) and pka and pKb

values. These acids binary (HX like HCl) or Oxo acid (HaXbCc like H2SO4); the strength

of binary acid increase with the size of X(HI>HBr>HCl>HF since I->Br->Cl->F- in size)

while the strength of oxoacid increased with electro negativity of X, No.of oxygen atoms

and the oxidation state of central element X.

Review Questions 1. In the acid-base forward reaction of HF + H2O →H3O++ F–, identify the acid, base and

their conjugates.

A. Acid = ______________ C. Conjugate Acid = ______________

B. Base = ______________ D. Conjugate Base = ______________

2. Which acid pairs is the stronger acid? Explain your answer in clear way.

A. HClO and HClO2 C. HBrO3 or HBrO4

B. H2SO4 and H3PO4 D. HOCl or HOBr

3. Write at least two differences between hard acid and soft acid.

4. How are acids and bases defined in terms of (i) Arrhenius concept and (ii) Bronsted

Lowry concept? Give suitable examples.

5. Explain HSAB principle.

4. Periodic Table

General Properties of Main Group Element

At the end of this session, students will be able to:

Explain the general properties of s-block elements

Describe s-block elements

State the periodic and group properties of main group elements

Familiar with periodicity in main group elements

Neighboring effect of elements in main group elements




1. What is the main difference among p-block element, d-block elements and inner

transition elements? What are their physical and chemical properties?

2. What is periodic table? How elements can arrange in periodic table?

4.1. Introduction In this section we will study the atomic structure and correlate chemical and physical

properties of the elements. First we will focus the physical properties such as atomic

and ionic size, ionization energy, and electron affinity, electro negativity and next we

will precede the periodicity of chemicals properties such as metallic and non metallic

nature, oxidation state, and bond character.

4.2. Periodicity of physical properties In 1869 and 1870 respectively, Dmitri Mendeleev and Lothar Meyer stated that the

properties of the elements can be represented in periodic table as the functions of their

atomic weights, and set out their ideas in the form of a periodic table. As new

elements have been discovered, the original form of the periodic table has been

extensively modified, and it is now recognized that periodicity is a consequence of the

variation in ground state electronic configurations.

A modern periodic table emphasizes the blocks of 2, 6, 10 and 14 elements which

result from the filling of the s, p, d and f atomic orbital’s respectively. An exception is

He, which, for reasons of its chemistry, is placed in a group with Ne, Ar, Kr, Xe and

Rn.



The modern periodic table in which the elements are arranged in numerical

orders according to the number of protons, they possess. The division into

groups places elements with the same number of valence electrons into

vertical columns within the table. The vertical groups of elements are called

triads. Rows in the periodic table are called periods. The first period contains

H and He, but the row from Li to Ne is sometimes referred to as the first

period.

4.3. Periodic Variation in Physical Properties

4.3.1. Atomic Radii As we move across the periodic table, atoms become smaller due to increasing effective

nuclear charges. The effective nuclear charge, Zeff experienced by an electron in an

outer shell is less than the actual nuclear charge, Z. This is because the attraction of

outer-shell electrons by the nucleus is partly counter balanced by the repulsion of

these outer-shell electrons by electrons in inner shells. This concept of a screening, or

shielding, effect helps us understand many periodic trends in atomic properties.

Consider an atom of lithium; it has two electrons in a filled shell, 1s2, and one electron

in the 2s orbital, 2s1. The electron in the 2s orbital is fairly effectively screened from

the nucleus by the two electrons in the filled 1s orbital, so the 2s electron does not



―feel‖ the full 3 charge of the nucleus. The effective nuclear charge, Zeff, experienced

by the electron in the 2s orbital: ZZeff where Z=atomic number and ð is core

electrons. Example;11Na ,Zeff=11-10 = +1,12Mg, Zeff=12-10= +2,13Al, Zeff=13-10=

+3.

In general, as we move from left to right across a period in the periodic table, atomic

radii of representative elements decrease as a proton is added to the nucleus and an

electron is added to a particular shell. Down the group atomic size increased since

there are additional shells.

4.3.2. Ionic Radii Ionic Radii of metals always greater than the corresponding cations because they lost

electrons from shells which reduced electron-electron repulsions. Example;Mg (1.6

Ǻ) > Mg2+(0.85 Ǻ).but the radii of negatively charged ions (anions) are always larger

than the neutral atoms from which they are formed because there is addition of extra

electron to the shell which increased electron-electron. Example: Cl- (1.67 Ǻ) > Cl (1.1

Ǻ). Both sizes of cations and anions decrease from left to right across a period and

increase going down a group.

Within an isoelectronic series, radii decrease with increasing atomic number

because of increasing nuclear charge. Isoelectronic ions: ions with the same

number of core electrons. Na+, Mg+2, Al+3, F-, O-2, N-3 all contain 10 electrons; all have

the same electron configuration as Ne but, in terms of size, N-3>O-2>F-

>Na+>Mg+2>Al+3 ;the ion with the greater number of protons in an isoelectronic series

will be the smallest due to the greater nuclear charge pulling the electrons in closer.

Arrange the following ions in order of increasing ionic radii: (a) Ca2+, K+, Al3+ +, (b) Se2-,

Br-, Te2- Answers: a) Al3+< Ca2+ <K , b)Br-< Se2- < Te2-

4.3.3. Ionization Energy

Ionization (potential) Energy: the minimum amount of energy that required for

removing an electron from ground state (loosely bounded) atom in gaseous state. It

measure how outermost electrons held by the nucleus.

X(g)+ IE X+(g)+e-

The first ionization energy (IE1), also called first ionization potential, is the minimum

amount of energy required to remove the most loosely bound electron from an isolated

gaseous atom to form an ion with a 1+charge. Elements with low ionization energies

(IE) lose electrons easily to form cations (they are highly metallic).



In general ionization energy decrease down the group since Zeff remains constant; the

size of atoms increase which results the outermost electrons bound loosely. In general

ionization energy increase across the period since Zeff increase by one; the size of

atoms decrease which results the outermost electrons bound tightly with the nucleus.

The first ionization energies of the Group IIA elements (Be, Mg, Ca, Sr, Ba) are also

significantly higher than those of the Group IA elements in the same periods. This is

because the Group IIA elements have higher Zeff values and smaller atomic radii.

Thus, their outermost electrons are held more tightly than those of the neighboring IA

metals. It is harder to remove an electron from a pair in the filled outermost s orbitals

of the Group IIA elements than to remove the single electron from the half-filled

outermost s orbitals of the Group IA elements.

Exceptions occur at Groups IIIA and VIA. The first ionization energies for the Group

IIIA elements (B, Al, Ga, In, Tl) are exceptions to the general horizontal trends. They

have lower than those of the IIA elements in the same periods because the IIIA

elements have only a single electron in their outermost p orbitals. Less energy is

required to remove the first p electron than the second s electron from the outermost

shell, because the p orbital is at a higher energy (less stable) than an s orbital within

the same shell (n value).

The general left-to-right increase in IE1 for each period is interrupted by a dip between

Groups VA (N, P, As, Sb, Bi) and VIA elements (O, S, Se, Te, Po). Presumably, this

behavior is because the fourth np electron in the Group VIA elements is paired with

another electron in the same orbital, so it experiences greater repulsion than it would

in an orbital by itself. This increases repulsion apparently outweighs the effect of Zeff,

so the fourth np electron in an outer shell (Group VIA elements) is somewhat easier to

remove (lower ionization energy) than is the third np electron in an outer shell (Group

VA elements). The first ionization energies generally increase from left to right across

the periodic table. The order of ionization energies decreased as 3rd > 2nd >1st because

it is too difficult to remove an electron from cation specie.



4.3.4. Electron Affinity (EA) The electron affinity (EA) of an element may be defined as the amount of energy

released when an electron is added to an isolated gaseous atom to form a negative

charge. X(g)+ e X-(g)+EA

Electron Affinity (EA): Energy is always required to bring a negative charge (electron)

closer to another negative charge (anion). So the addition of a second electron to anion

to form an ion with a -2charge is always endothermic. Thus, electron affinities of

anions are always positive. X-(g) + e → X2-(g)+EA

Most elements have no affinity for an additional electron and thus have an electron

affinity (EA) equal to zero. We can represent the electron affinities of helium and

chlorine as:

Halodens have ns2-np5 electron cofiguration which are most electron affinities to form

noble gas cofiguration. Electron affinity involves the addition of an electron to a

neutral gaseous atom. X (g) +e- →X-(g) + EA; In general EA of the element:

increase except noble gas

decrease

EA

4.3.5. Electronegativity (EN) The electro-negativity (EN) of an element is a measure of the relative tendency of an

atom to attract electrons towards it when it is chemically combined with another atom.

Elements with high electro-negativities (nonmetals) often gain electrons to form

anions. Example; F (4.0) > O(3.5) > Cl=N(3.0). Elements with low electro-negativities

(metals) often lose electrons to form cations. For the representative elements, electro-

negativities usually increase from left to right across periods except noble gas and

decrease from top to bottom within groups.

increase except noble gas

decrease

EA



Arrange the following elements in order of increasing electro-negativity? B, Na, F, O

ans, Na< B< O< F

4.3.6. Melting and boiling points The melting and boiling points of halogens increase with increase in atomic number

down the group. Becuase the forces existing between these molecules are weak

Van der Waals forces, which increase down the group. This is also clear from

the change of state from fluorine to iodine. At room temperature, fluorine and

chlorine are gases; bromine is a liquid while iodine and astatine are solids.

4.3.7. Metallic character Metallic character increases from top to bottom and decreases from left to right with

respect to position in the periodic table. Nonmetallic character decreases from top to

bottom and increases from left to right in the periodic table.

4.3.8. Horizontal, Vertical, and Diagonal Relationships On moving from left to right across a period, the size of the atom decreases because

nuclear charge is increased. Thus, the orbital electrons are more tightly held, and the

ionization energy increases. Hence, metallic character of the element decreased. On

descending a group in the periodic table, the element all have the same number of

outer electrons and the same valency, but the size increases. Thus, the ionization

energy decreases and the metallic character increases. On moving diagonally across

the periodic table the element shows certain similarities. These are usually weaker

than the similarities with in a group, but are quite pronounced in the following pairs

of elements.

Dear students! The detail properties, electron configuration, ocurance, important of main

group elements (from group IA to VIIIA) are left for group assignment.

+

L i B e

B C N a M g A l



Unit Summary In modern periodic table, elements are arranged as the function of their atomic

number. Main group elements are elements contain s-block metals and p-block metal,

metalloid and nonmetals. This periodic table constructed with the columns (family)

and the horizontal row (periods). Periodicity is the variation of physical properties in

periodic table such as atomic size, electronegativity, ionization energy, metallic

character and others as summarized below.

Review Questions

1. Define the followings.

A. Period

B. Group

C. Representative elements

3. How to determine group numbers and periods of main group elements?



5. Chemistry of Hydrogen

At the end of this session, students will be able to:

Explain the general properties of Hydrogen element

Describe the reactivity of hydrogen

Familiar with the position of hydrogen periodic table

Important of hydrogen and hydrogen compounds

Brainstorming Question

What does mean hydrogen? How to prepare hydrogen?

5.1. History of Hydrogen Hydrogen was first isolated and shown to be a discrete element by Henry Cavendish in

1766. Before that, Robert Boyle and Paracelsus both used reactions of iron and acids to

produce hydrogen gas. Hydrogen comes from Greek meaning ―water producer‖ (―hydro”

=water and “gennao”=to make). Antoine Lavoisier gave hydrogen its name because it

produced water when ignited in air. Hydrogen is one of the most important elements in the

world. It is all around us. It is a component of water (H2O), fats, petroleum, table sugar

(C6H12O6), ammonia (NH3), and hydrogen peroxide (H2O2)—things essential to life, as we know

it.

5.2. Position of Hydrogen in Periodic Table The placement of hydrogen element in periodic table is a difficult task for chemists. This is due

to the fact that hydrogen shows both metallic and nonmetallic property. In light this, different

chemists placed hydrogen element in different groups.

Most scientist placed hydrogen element in group IA since both hydrogen and alkali metals

react with nonmetals (halogen) to have +1 oxidation state. However, it varies greatly from the

alkali metals as it forms hydride ion (H-) but, alkali metals never been. Besides, hydrogen has

high ionization energy (1312 kJ/mol), while lithium (the alkali metal with the highest ionization

energy) has an ionization energy of 520 kJ/mol. Because hydrogen is a nonmetal and forms H-

(hydride anions), it is sometimes placed above the halogens in the periodic table.

Some scientist also placed hydrogen in group VIIA since both hydrogen and halogen react with

alkali metals to full fill their electron deficiency. Hydrogen also forms H2 dihydrogen like

halogens. However, hydrogen is very different from the halogens. Hydrogen has a much

smaller electron affinity than the halogens. H2 dihydrogen or molecular hydrogen is non-

polar with two electrons. However, H2 has very strong intramolecular forces; H2 reactions are

generally slow at room temperature due to strong H—H bond. H2 is easily activated by heat,



irradiation, or catalysis. Activated hydrogen gas reacts very quickly and exothermically with

many substances. Hydrogen also has an ability to form covalent bonds with a large variety of

substances. Because it makes strong O—H bonds, it is also considered a good reducing agent

for metal oxides. Example: CuO(s) + H2(g) → Cu(s) + H2O(g) H2(g) passes over CuO(s) to reduce

the Cu2+ to Cu(s), while getting oxidized itself.

A few scientist placed hydrogen in group IVA since both hydrogen and carbon family form a

covalent bond by sharing of their electrons.

After a long time argument, in modern periodic table hydrogen is placed above group IA in the

periodic table because it has ns1 electron configuration like the alkali metals.

5.3. Isotopes of Hydrogen

Protium Deuterium Tritium

Protium (1H) is the most common isotope, consisting of 99.98% of naturally occurring

hydrogen. It is a nucleus containing a single proton. Deuterium (2H) is another isotope

containing a proton and neutron, consisting of only .0156% of the naturally occurring

hydrogen. Commonly indicated with symbol D, D20 is called heavy water, which has a higher

density, melting point, and boiling point than regular water. Replacing protium with

deuterium (called deuteration) has important implications for the rate of reaction called kinetic-

isotope effect. Tritium (3H) is a radioactive isotope with a 12.3-year half-life, which is

continuously formed in the upper atmosphere due to cosmic rays. It is can also be made in a

lab from Lithium-6 in a nuclear reactor. Tritium is also used in hydrogen bombs.

5.4. Property and Occurrence of Hydrogen Hydrogen is a nonmetal. At room temperature, hydrogen is colorless and odorless diatomic

gas. Hydrogen is the lightest and most abundant element in the universe. About 70%- 75% of

the universe is composed of hydrogen by mass. All stars are essentially large masses of

hydrogen gas that produce enormous amounts of energy through the fusion of hydrogen atoms

at their dense cores. In smaller stars, hydrogen atoms collided and fused with helium and

other light elements like nitrogen and carbon (essential for life). In the larger stars, fusion

produces the lighter and heavier elements like calcium, oxygen, and silicon.

On Earth, hydrogen is mostly found in association with oxygen; it’s most abundant form being

water (H2O). Hydrogen is only 0.9% by mass and 15% by volume abundant on the earth,

despite water covering about 70% of the planet. Because hydrogen is so light, there is only 0.5



ppm (parts per million) in the atmosphere, which is a good thing considering it is EXTREMELY

flammable.

5.5. Reactions of Hydrogen

Hydrogen wanting to give up its single electron causes it to act like an alkali metal: H(g)→

H+(g) + e-

A half-filled valence shell with one electron- also causes hydrogen to act like a halogen

because it wants to gain a Noble gas electron configuration by adding an e-: H(g) + e- → H-

(g)

Reactions with Active Metals: Hydrogen accepts electron from an active metal to form

ionic hydrides like LiH. By forming an ion with -1 charge, the hydrogen behaves like a

halogen.

2M(s)+H2(g) 2MH(s) ,where M=group IA

metals Example: 2K(s)+H2(g) → 2KH(s) 2K(s)+Cl2(g) → 2KCl(s)

Reactions with nonmetals: Unlike metals forming ionic bonds with nonmetals, hydrogen

forms polar covalent bonds (like the active metals that form ionic bonds with nonmetals),

but hydrogen is much less electropositive than the active metals, and forms covalent

bonds.

Hydrogen + Halogen → Hydrogen Haliden(H2(g)+ Cl2(g) → HCl(g))

Hydrogen + Oxygen → Water (H2(g) + O2(g) → H2O(g))

Reactions with Transition Metals: Reactions of hydrogen with transition metals (Group

3-12) form metallic hydrides. There is no fixed ratio of hydrogen atom to metal because

the hydrogen atoms fill holes between metal atoms in the crystalline structure.

5.6. Uses and Application of Hydrogen Hydrogen is very important to this world. About 70% of the hydrogen produced is used in the

Haber process, which is a process of fixing nitrogen gas into ammonia (a usable form by

plants). Without the Haber process, we would not be able to grow the huge amounts of crops

we grow today. Hydrogen is also used for the hydrogenation of oils. Hydrogenation entails

replacing double bonds in oils by hydrogen, converting the double bonds into single

bonds. This transformation of unsaturated fats to saturated fats drastically increases the shelf

life of many foods. However, an increased consumption of saturated fats has been linked to

greater visible for heart disease, high cholesterol, and certain types of cancer. Because

hydrogen is a good reducing agent, it is used to produce metals like iron, copper, nickel, and

cobalt from their ores.

Liquid hydrogen (combined with liquid oxygen) is a major component of rocket fuel (as

mentioned above combination of hydrogen and oxygen relapses a huge amount of energy).



However, the use of hydrogen for this purpose was largely discontinued around World War II

after the explosion of The Hindenburg. The Hindenburg prompted greater use of inert helium,

rather than flammable hydrogen for air travel. Recently, due to the fear of fossil fuels running

out, extensive research is being done on hydrogen as a source of energy. Because of their

moderately high energy densities liquid hydrogen and compressed hydrogen gas are possible

fuels for the future. A huge advantage in using them is that their combustion only produces

water (it burns ―clean‖). However, it is very costly, and not economically feasible with current

technology.

Combustion of fuel produces energy that can be converted into electrical energy when energy

in the steam turns a turbine to drive a generator. However, this is not very efficient because a

great deal of energy is lost as heat. The production of electricity using voltaic cell can yield

more electricity (a form of usable energy). Voltaic cells that transform chemical energy in fuels

(like H2 and CH4) are called fuel cells. These are not self-contained and so are not considered

batteries. The hydrogen cell is a type of fuel cell involving the reaction between H2(g) with O2(g)

to form liquid water; this cell is twice as efficient as the best internal combustion engine. In

the cell (in basic conditions), the oxygen is reduced at the cathode, while the hydrogen is

oxidized at the anode.

Reduction: O2(g)+2H2O(l)+4e- → 4OH-(aq)

Oxidation: H2(g) + 2OH-(aq) → 2H2O(l) + 2e-

Overall: 2H2(g) + O2(g) → 2H2O(l)

Ecell= Reduction- Oxidation= E O2/OH- - E H2O/H2 = .401V – (-.828V) = +1.23

5.7. Preparation of Hydrogen Hydrogen gas can be prepared by reacting with dilute strong acid like hydrochloric acids

with an active metal. The metal becomes oxides, while the H+ (from the acid) gests reduced

to hydrogen gas. This method is only practical for producing small amounts of hydrogen in

the lab, but is much too costly for industrial production: Zn(s) + 2H+(aq)→ Zn2+(aq) + H2(g)

The purest form of H2(g) can prepared from electrolysis of H2O(l), the most common

hydrogen compound on this plant. This method is also not commercially viable because it

requires a huge amount of energy (about 572 kJ): 2H2O(l) → 2H2(g) + O2(g) ΔH°=+572

kJ

H2O is the most abundant form of hydrogen on the planet, so it seems logical to try to

extract hydrogen from water without electrolysis of water. To do so, we must reduce the

hydrogen with +1 oxidation state to hydrogen with 0 oxidation state (in hydrogen gas). Three

commonly used reducing agents are carbon (in coke or coal), carbon monoxide, and



methane (Reforming of Methane). These react with water vapor form H2(g). These three

methods are industrially feasible (cost effective) methods of producing H2 (g).

C(s) + 2H2O(g) → CO(g) + H2(g)

CO(g) + 2H2O(g) → CO2 + H2(g)

CH4(g) + H2O(g) → CO(g) + 3H2(g)

Unit Summary Many chemists have argued for the position of hydrogen in periodic table; most of

them placed in IA metals since both alkali metal and H have one valence electron;

both react with halogen to form (NaCl and HCl); but unlike IA, H has high ionization

energy, H can share electron to form covalent bond). A few chemists place H in VIIA

because both H and halogens are one electron deficient and they react with Metals

(e.g. NaH and NaCl); but halogen have high electronegative and affinity. Few chemists

place H in group IVA since both carbon and H can form a covalent bond by sharing

electron. In light these, now a day hydrogen place above the IA elements because H is

more similar with alkali metals.

Review questions 1. What are the similarity and differences between hydrogen and halogens,

alkali metals or carbon family?

2. Write some physical and chemical properties of Hydrogen.

3. How can prepare hydrogen? Write some preparation methods.

4. List some use of hydrogen and its compounds.

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Inorganic Chemistry- I 1. Atomic Theory and Atomic Structure

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