mst,module 1, notes

7/29/2019 MST,Module 1, Notes

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4.5.1:- Module 1

4.5.1.1:- Introduction to Materials:

Scope of Material Science & Technology and its Importance in Chemical Engineering:-

Materials science is an interdisciplinary field applying the properties of matter to various

areas ofscience and engineering. This scientific field investigates the relationship between the

structure of materials at atomic or molecular scales and their macroscopic properties. It

incorporates elements ofapplied physicsandchemistry. With significant media attention focusedon nanoscience and nanotechnologyin recent years. It is also an important part offorensicengineering and failure analysis.

Materials science also deals with fundamental properties and characteristics of materials.

The basis of materials science involves relating the desiredproperties and relative performance

of a material in a certain application to the structure of the atoms and phases in that material

through characterization. The major determinants of the structure of a material and thus of itsproperties are its constituent chemical elements and the way in which it has been processed into

its final form. These characteristics, taken together and related through the laws of

thermodynamics, govern a materials microstructure, and thus its properties.Materials science encompasses various classes of materials, each of which may constitute a

separate field. There are several ways to classify materials. For instance by the type of bonding

between the atoms. The traditional groups are ceramics, metals and polymers based on atomic

structure and chemical composition. New materials have resulted in more classes. One way ofclassifying materials is:

Biomaterials

Carbon

Ceramics

Composite materials

Glass

Metals

Nanomaterials

Polymers

Refractory

Semiconductors

Thin Films

Functionally Graded Materials.

Radical materials advances can drive the creation of new products or even new industries, butstable industries also employ materials scientists to make incremental improvements and

troubleshoot issues with currently used materials. Industrial applications of materials science

include materials design, cost-benefit tradeoffs in industrial production of materials, processingtechniques (casting,rolling,welding, ion implantation, crystal growth, thin-film deposition,

sintering, glassblowing, etc.), and analytical techniques (characterization techniques such as

Electron microscopy, X-ray diffraction,Calorimetry,Nuclear microscopy, Rutherfordbackscattering,Neutron diffraction, small-angle X-ray scattering, etc.).

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Besides material characterization, the material scientist/engineer also deals with the extraction of

materials and their conversion into useful forms.

Thus ingot casting, foundry techniques, blast furnace extraction, and electrolytic extraction areall part of the required knowledge of a metallurgist/engineer. Often the presence, absence or

variation of minute quantities of secondary elements and compounds in a bulk material will have

a great impact on the final properties of the materials produced, for instance, steels are classified

based on 1/10 and 1/100 weight percentages of the carbon and other alloying elements theycontain. Thus, the extraction and purification techniques employed in the extraction of iron in the

blast furnace will have an impact of the quality of steel that may be produced.

4.5.1.2:- Atomic Structure & Chemical Bonding:-

Structure of the Atom: The name atom comes from the Greekword atomos" which meansimpossible to cut, or indivisible, something that cannot be divided further. The concept of an

atom as an indivisible component of matter was first proposed by early Indian and Greek

philosophers.

Basic Atomic Structure

Atoms are made up of three main particles. These

are the neutron, electron, and proton. These main

parts are each made up of smaller particles. This picture shows the placement of each of the main

particles.

The Proton is positively charged and is in the nucleus.

The Neutron has no electrical charge and is in the nucleus also.

The electron moves around the nucleus and it has a negative charge.

The number of protons in an atom determines the type of atom it will be.

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The proton and the neutron are very large and make up most of the MASS of the atom. The

electrons have very little mass, but since they move around the nucleus they take up most of the

SPACE in the atom.
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A generic atomic planetary model, or the Rutherford model:-

A Bohr model of the hydrogen atom, showing an electron jumping between fixed orbits and

emitting a photon of energy with a specific frequency:-

The behavior of protons, neutrons and electrons in electric fields:-

If a beam of each of these particles is passed between two electrically charged plates - onepositive and one negative-

Protons are positively charged and so would be deflected on a curving path towards the negative

plate.

Electrons are negatively charged and so would be deflected on a curving path towards the

positive plate.

Neutrons don't have a charge, and so would continue on in a straight line.

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Nucleus:

The nucleus is at the centre of the atom and contains the protons and neutrons. Protons and

neutrons are collectively known as nucleons.

Virtually all the mass of the atom is concentrated in the nucleus, because the electrons weigh so

little.

Working out the numbers of protons and neutrons

No of protons = ATOMIC NUMBER of the atom

The atomic number is also given the more descriptive name ofproton number.

No of protons + no of neutrons = MASS NUMBER of the atom

The mass number is also called the nucleon number.

This information can be given simply in the form:

How many protons and neutrons has this atom got?

The atomic number counts the number of protons (9); the mass number counts protons +

neutrons (19). If there are 9 protons, there must be 10 neutrons for the total to add up to 19.

The atomic number is tied to the position of the element in the Periodic Table and therefore thenumber of protons defines what sort of element you are talking about. So if an atom has 8

protons (atomic number = 8), it must be oxygen. If an atom has 12 protons (atomic number =

12), it must be magnesium.

Similarly, every chlorine atom (atomic number = 17) has 17 protons; every uranium atom(atomic number = 92) has 92 protons.

The atom is a basic unit ofmatterthat consists of a dense central nucleus surrounded by a cloud

ofnegatively chargedelectrons. Theatomic nucleus contains a mix of positively chargedprotons

and electrically neutral neutrons(except in the case ofhydrogen-1, which is the only stablenuclidewith no neutrons). The electrons of an atom are bound to the nucleus by the

electromagnetic force. Likewise, a group of atoms can remain bound to each other by chemical

bonds based on the same force, forming a molecule. An atom containing an equal number of

protons and electrons is electrically neutral, otherwise it is positively or negatively charged and isknown as an ion. An atom is classified according to the number of protons and neutrons in its

nucleus: the number of protonsdetermines thechemical element, and the number of neutrons

determines the isotope of the element.
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occurring elements have also been produced in laboratories.

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Ionization Potential or Ionization Energy:Amount ofenergy required to remove anelectron from an isolated atom ormolecule. There is an

ionization potential for each successive electron removed, though that associated with removingthe first (most loosely held) electron is most commonly used.

The ionization potential of an element is a measure of its ability to enter into chemical reactionsrequiring ion formation or donation of electrons and is related to the nature of the chemical

bonding in the compounds formed by elements.

Thus ionization potential can be defined as the energy per unit charge needed to remove anelectron from a given kind of atom or molecule to an infinite distance; usually expressed in volts.

Ionization potential is the potential difference through which a bound electron must be raised tofree it from the atom or molecule to which it is attached. In particular, the ionization potential is

the difference in potential between the initial state, in which the electron is bound, and the final

state, in which it is at rest at infinity.

The ionization potential for the removal of an electron from a neutral atom other than hydrogen

is more correctly designated as the first ionization potential. The potential associated with the

removal of a second electron from a singly ionized atom or molecule is then the secondionization potential, and so on.

Electron affinity:Definition
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The electron affinity of anatom ormolecule is definedas the amount of energy releasedwhen

an electron is added to a neutral atom or molecule to form a negative ion.

X + e X

This property is measured for atoms and molecules in the gaseous state only, since in the solid or

liquid states their energy levels would be changed by contact with other atoms or molecules.

A list of the electron affinities was used by Robert S. Mulliken to develop an electronegativityscale for atoms, equal to the average of the electron affinity and ionization potential.

Other theoretical concepts that use electron affinity include electronic chemical potential and

chemical hardness.

Another example, a molecule or atom that has a more positive value of electron affinity than

another is often called anelectron acceptorand the less positive an electron donor. Together they

may undergo charge-transferreactions.

In solids, the electron affinity is the energy difference between thevacuum energy and the

conduction band minimum.

To use electron affinities properly, it is essential to keep track of sign.

For any reaction that releases energy, the change in energy, E, has a negative value and the

reaction is called anexothermic process.

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Electron capture for almost all non-noble gas atoms involves the release of energyand thus are

exothermic. The positive values that are listed in tables ofEea are amounts or magnitudes. It isthe word, releasedwithin the definition energy releasedthat supplies the negative sign.

Confusion arises in mistakingEea for a change in energy, E, in which case the positive values

listed in tables would be for an endo- not exo-thermic process. The relation between the two is

Eea = - E(attach).However, if the value assigned toEea is negative, the negative sign implies a reversal of direction,

and energy is requiredto attach an electron. In this case, the electron capture is an endothermic

process and the relationship,Eea = - E(attach) is still valid. Negative values typically arise forthe capture of a second electron, but also for the nitrogen atom.

The usual expression for calculatingEea when an electron is attached is

Eea = (Einitial Efinal)attach = - E(attach)

This expression does follow the convention X = X(final) X(initial) since E = - (E(final)

E(initial)) = E(initial) E(final).

Equivalently, electron affinity can also be defined as the amount of energy requiredto detach anelectron from a singly chargednegative ion,i.e. the energy change for the process

X X + e

If the same table is employed for the forward and reverse reactions, without switching signs, care

must be taken to apply the correct definition to the corresponding direction, attachment-(release)

or detachment-(require). Since almost all detachments (require +) an amount of energy listed on

the table, those detachment reactions are endothermic, or E(detach) > 0.
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Eea = (Efinal Einitial)detach = E(detach) = - E(attach)

Electronegativity:Definition

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

The Pauling scale is the most commonly used. Fluorine (the most electronegative element) isassigned a value of 4.0, and values range down to caesium and francium which are the leastelectronegative at 0.7.

What happens if two atoms of equal electronegativity bond together?

Consider a bond between two atoms, A and B. Each atom may be forming other bonds as well as

the one shown but these are irrelevant to the argument.

If the atoms are equally electronegative, both have the same tendency to attract the bonding pair

of electrons, and so it will be found on average half way between the two atoms. To get a bondlike this, A and B would usually have to be the same atom. You will find this sort of bond in, for

example, H2 or Cl2 molecules.

This sort of bond could be thought of as being a pure covalent bond where the electrons are

shared evenly between the two atoms.

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What happens if B is slightly more electronegative than A?

B will attract the electron pair rather more than A does.

That means that the B end of the bond has more than its fair share of electron density and so

becomes slightly negative. At the same time, the A end (rather short of electrons) becomes

slightly positive. In the diagram, " " (read as "delta") means "slightly" - so + means "slightly

positive".

This is described as aPolar bond. A polar bond is a covalent bond in which there is a separationof charge between one end and the other - in other words in which one end is slightly positive

and the other slightly negative. Examples include most covalent bonds. The hydrogen-chlorinebond in HCl or the hydrogen-oxygen bonds in water are typical.

What happens if B is a lot more electronegative than A?

In this case, the electron pair is dragged right over to B's end of the bond. To all intents andpurposes, A has lost control of its electron, and B has complete control over both electrons. Ions

have been formed.


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Time-dependent Schrdinger equation(general)

where (psi)is thewave function of the quantum system, iis the imaginary unit,is the

reduced Planck constant, and is the Hamiltonianoperator, which characterizes the total energy

of any given wave-function and takes different forms depending on the situation.

A wave function which satisfies the non-relativistic Schrdinger equation with V=0. In other

words, this corresponds to a particle traveling freely through empty space. The real part of the

wave functionis plotted here.

The most famous example is thenon-relativisticSchrdinger equation for a single particlemoving in an electric field (but not a magnetic field, see thePauli equation):

Time-dependent Schrdinger equation(single non-relativistic particle)

where, m is the particle's mass, Vis itspotential energy,2 is the Laplacian, and is the wave-function (more precisely, in this context, it is called the

"position-space wave-function").

In plain language, it means "total energy equals kinetic energy pluspotential energy".

[In mathematicsthe Laplace operator orLaplacian is a differential operatorgiven by the divergenceof thegradient of afunction onEuclidean space. It is usually denoted by the symbols

,2 or . The Laplacian occurs indifferential equations that describe many physical phenomena, such as electricandgravitational

potentials, the diffusion equationforheatand fluid flow,wave propagation, andquantum mechanics.]

The term "Schrdinger equation"can refer to both the general equation (first box above), or the

specific non-relativistic version (second box above and variations thereof). The general equationis indeed quite general, used throughout quantum mechanics.

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To apply the Schrdinger equation, the Hamiltonian operator is set up for the system, accounting

for the kinetic and potential energy of the particles constituting the system, then inserted into the

Schrdinger equation. The resulting partial differential equation is solved for the wave-function,

which contains information about the system.
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Time-independent equation

Each of these three rows is a wave-function which satisfies the time-dependent Schrdinger

equation for a harmonic oscillator. Left: The real part (blue) and imaginary part (red) of the

wave-function. Right: Theprobability distribution of finding the particle with this wave-function

at a given position. The top two rows are examples ofstationary states, which correspond to

standing waves. The bottom row an example of a state which is nota stationary state. The right

column illustrates why stationary states are called "stationary".

The time-dependent Schrdinger equation predicts that wave-functions can form standing

waves, called stationary states (also called "orbitals", as in atomic orbitals ormolecular orbitals).These states are important in their own right, and moreover if the stationary states are classified

and understood, then it becomes easier to solve the time-dependent Schrdinger equation forany

state. The time-independent Schrdinger equation is the equation describing stationary states. (Itis only used when the Hamiltonian itself is not dependent on time.)

Time-independent Schrdinger equation (general)

In words, the equation states:When the Hamiltonian operator acts on the wavefunction , the result might be

proportional to the same wavefunction . If it is, then is a stationary state, and the

proportionality constant, E, is the energy of the state .

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Chemical Bonding:
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Some of the most essential facts on Atomic Structure and Chemical Bonding are :

1. Elements are made up of atoms.

2. Each atom has a nucleus situated at the center. It contains positively charged particles called

protons, and neutral particles called neutrons.

3. Electrons are negatively charged particles which move around the nucleus in definite circular

paths called orbits, shells or energy levels.4. The mass number of an element is equal to the sum of the number of protons and number of

neutrons in its nucleus.

5. Number of protons equals the number of electrons in an atom; therefore, an atom is

electrically neutral.

6. Atomic number is the number of protons of an atom.

7. Isotopes are atoms of the same element having different mass numbers.

8. The distribution of electrons in various shells or energy levels in an atom is called the

electronic configuration of that atom.

9. According to Bohr and Bury, the maximum number of electrons that can be accommodated in

any energy level of an atom is given by the formula 2n

2

, where n represents the number ofthe energy level.

10. In order to exist independently by itself an atom must have eight electrons and if in its

outermost shell two electrons then there is only one shell. This is the octet rule.

11. Atoms try to attain stable configuration (completing their outermost shell) either by losing,

gaining or sharing electrons.

12. The force of attraction that holds atoms together in a molecule is known as a chemical bond.

13. A bond between an anion and cation is called an ionic bond. Cations give electrons to the

anions.

14. A covalent bond is a bond in which both the reacting atoms are short of electrons. Thus, they

attain stable electronic configuration by sharing electrons.15. Coordinate bond is a covalent bond in which the shared pair of electrons is contributed by

only one of the two atoms.

A chemical bond is an attraction between atoms that allows the formation ofchemicalsubstancesthat contain two or more atoms.

The bond is caused by the electrostatic force of attraction between opposite charges, eitherbetween electronsand nuclei, or as the result of adipole attraction.
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Main types of chemical bonds:

A chemical bond is an attraction between atoms. This attraction may be seen as the result of

different behaviors of the outermost electrons of atoms. Although all of these behaviors

merge into each other seamlessly in various bonding situations so that there is no clear line

to be drawn between them, nevertheless behaviors of atoms become so qualitatively

different as the character of the bond changes quantitatively, that it remains useful andcustomary to differentiate between the bonds that cause these different properties of

condensed matter.

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In covalent bond, one or more electrons (often a pair of electrons) are drawn into the spacebetween the two atomic nuclei. Here the negatively charged electrons are attracted to the positive

charges ofboth nuclei, instead of just their own. This overcomes the repulsion between the two

positively charged nuclei of the two atoms, and so this overwhelming attraction holds the two

nuclei in a fixed configuration of equilibrium, even though they will still vibrate at equilibriumposition. Thus, covalent bonding involves sharing of electrons in which the positively charged

nuclei of two or more atoms simultaneously attract the negatively charged electrons that arebeing shared between them.

In a polar covalent bond, one or more electrons are unequally shared between two nuclei.Covalent bonds often result in the formation of small collections of better-connected atoms

called molecules, which in solids and liquids are bound to other molecules by forces that are

often much weaker than the covalent bonds that hold the molecules internally together. Suchweak intermolecular bonds give organic molecular substances, such as waxes and oils, their soft

bulk character, and their low melting points (in liquids, molecules must cease most structured or

oriented contact with each other).

When covalent bonds link long chains of atoms in large molecules, however (as in polymers

such as nylon).

Or when covalent bonds extend in networks though solids that are not composed of discretemolecules (such as diamond orquartz or the silicate minerals in many types of rock) then the

structures that result may be both strong and tough, at least in the direction oriented correctly

with networks of covalent bonds. Also, the melting points of such covalent polymers andnetworks increase greatly.

In an ionic bond, the bonding electron is not shared at all, but transferred. In this type of bond,

the outeratomic orbitalof one atom has a vacancy which allows addition of one or more

electrons. These newly added electrons potentially occupy a lower energy-state (effectivelycloser to more nuclear charge) than they experience in a different atom. Thus, one nucleus offers

a more tightly bound position to an electron than does another nucleus, with the result that one

atom may transfer an electron to the other. This transfer causes one atom to assume a net positive

charge, and the other to assume a net negative charge.The bondthen results from electrostatic attraction between atoms, and the atoms become positiveor negatively charged ions.

Ionic bonds may be seen as extreme examples of polarization in covalent bonds. Often, such

bonds have no particular orientation in space, since they result from equal electrostatic attraction

of each ion to all ions around them.

Ionic bonds are strong (and thus ionic substances require high temperatures to melt) but alsobrittle, since the forces between ions are short-range, and do not easily bridge cracks and
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fractures. This type of bond gives a characteristic physical character to crystals of classic mineral

salts, such as table salt.

A metallic bond. In this type of bonding, each atom in a metal donates one or more electrons to a

"sea" of electrons that reside between many metal atoms. In this sea, each electron is free (byvirtue of its wave nature) to be associated with a great many atoms at once. The bond results

because the metal atoms become somewhat positively charged due to loss of their electrons,

while the electrons remain attracted to many atoms, without being part of any given atom.Metallic bonding may be seen as an extreme example of delocalization of electrons over a large

system of covalent bonds, in which every atom participates.

This type of bonding is often very strong (resulting in the tensile strength of metals). However,

metallic bonds are more collective in nature than other types, and so they allow metal crystals to

12

more easily deform, because they are composed of atoms attracted to each other, but not in any

particularly-oriented ways. This results in the malleability of metals. The sea of electrons inmetallic bonds causes the characteristically good electrical and thermal conductivity of metals,

and also their "shiny" reflection of most frequencies of white light.

Strength of chemical bonds: It varies considerably; there are "strong bonds" such ascovalent orionic bondsand "weak bonds" such as dipoledipole interactions, the London

dispersion force and hydrogen bonding.

Since opposite charges attract via a simple electromagnetic force, the negatively chargedelectronsthat are orbiting the nucleus and the positively chargedprotonsin the nucleus attract

each other. Also, an electron positioned between two nuclei will be attracted to both of them.

Thus, the most stable configuration of nuclei and electrons is one in which the electrons spendmore time between nuclei, than anywhere else in space. These electrons cause the nuclei to be

attracted to each other, and this attraction results in the bond. However, this assembly cannot

collapse to a size dictated by the volumes of these individual particles. Due to the matter wave

nature of electrons and their smaller mass, they occupy a much larger amount of volumecompared with the nuclei, and this volume occupied by the electrons keeps the atomic nuclei

relatively far apart, as compared with the size of the nuclei themselves.

In general, strong chemical bonding is associated with the sharing or transfer of electrons

between the participating atoms. The atoms in molecules, crystals, metals and diatomic gasesindeed most of the physical environment around us are held together by chemical bonds,

which dictate the structure and the bulk properties of matter.

Bond length:In molecular geometry, bond length orbond distance is the average distancebetween nuclei of twobondedatomsin a molecule. It is atransferable property of a bondbetween atoms of fixed types, relatively independent of the rest of molecule.

It is related tobond order, when more electrons participate in bond formation the bond will getshorter. Bond length is also inversely related tobond strength and thebond dissociation energy,

as (all other things being equal) a stronger bond will be shorter. In a bond between two identicalatoms half the bond distance is equal to the covalent radius.
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Bond lengths are measured in the solid phase by means ofX-ray diffraction, or approximated in

the gas phase by microwave spectroscopy. A set of two atoms sharing a bond is unique going

from one molecule to the next. For example the carbon to hydrogen bond in methane is differentfrom that in methyl chloride. It is however possible to make generalizations when the general

structure is the same.

Ionic (Electrovalent) bonding: We have already discussed this type above ( page: ).A simple view of ionic bonding:

The importance of noble gas structures: A lot of importance is attached to the electronic

structures of noble gases like neon or argon which have eight electrons in their outer energylevels (or two in the case of helium). These noble gas structures are thought of as being in some

way a "desirable" thing for an atom to have.

Ionic bonding in Sodium chloride

Sodium (2,8,1) has 1 electron more than a stable noble gas structure (2,8). If it gave away that

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electron it would become more stable.

Chlorine (2,8,7) has 1 electron short of a stable noble gas structure (2,8,8). If it could gain anelectron from somewhere it too would become more stable.

The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both become more

stable.

The sodium has lost an electron, so it no longer has equal numbers of electrons and protons.Because it has one more proton than electron, it has a charge of 1+. If electrons are lost from an

atom, positive ions are formed.

Positive ion is sometimes called cation.

The chlorine has gained an electron, so it now has one more electron than proton. It therefore hasa charge of 1-. If electrons are gained by an atom, negative ions are formed.

A negative ion is sometimes called an anion.

The nature of the bond

The sodium ions and chloride ions are held together by the strong electrostatic attractions

between the positive and negative charges, hence it is electrovalent or ionic bond.The formula of Sodium chloride

You need one sodium atom to provide the extra electron for one chlorine atom, so they combine

together 1:1. The formula is therefore NaCl.

Some other examples of ionic bonding:

Magnesium oxide

Again, noble gas structures are formed, and the magnesium oxide is held together by very strong

attractions between the ions. The ionic bonding is stronger than in sodium chloride because this

time you have 2+ ions attracting 2- ions. The greater the charge, the greater the attraction.
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The formula of magnesium oxide is MgO.Potassium oxide

Again, noble gas structures are formed. It takes two potassiums to supply the electrons the

oxygen needs. The formula of potassium oxide is K2O.

Covalent bonding: We have already discussed this type above ( page: ).A simple view of covalent bonding

The importance of noble gas structures: In this type too, there is a lot of importance is attached

to the electronic structures of noble gases like neon or argon which have eight electrons in theirouter energy levels (or two in the case of helium).

14As in the case of ionic bonding noble gas structures are achieved by transferring electrons from

one atom to another, it is also possible for atoms to reach these stable structures by sharing

electrons to give covalent bonds.

Some simple covalent molecules:

Chlorine:

For example, two chlorine atoms could both achieve stable structures by sharing their single

unpaired electron as in the diagram.

[The fact that one chlorine has been drawn with electrons marked as crosses and the other as dots to show from

where all the electrons come. In reality there is no difference between them.]

The two chlorine atoms are said to be joined by a covalent bond. The reason that the two chlorineatoms stick together is that the shared pair of electrons is attracted to the nucleus of both chlorine

atoms.

Hydrogen:

Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of

helium. Once again, the covalent bond holds the two atoms together because the pair of electrons

is attracted to both nuclei.

Hydrogen chloride:


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The hydrogen has a helium structure, and the chlorine an argon structure.

Some other cases:

Now, only thing which must be changed is the over-reliance on the concept of noble gas

structures. Most of the simple molecules you draw do in fact have all their atoms with noble gas

structures.

For example:

15

Even with a more complicated molecule like PCl3, theres no problem. In this case, only the outer

electrons are shown for simplicity. Each atom in this structure has inner layers of electrons of

2,8. Again, everything present has a noble gas structure.

Metallic bonding:We have already discussed this type above ( page: ).Metals tend to have high melting points and boiling points suggesting strong bonds between the

atoms.Metallic bonding in sodium:

A metal like sodium (melting point 97.8C) melts at a considerably higher temperature than theelement (neon) which precedes it in the Periodic Table.

Sodium has the electronic structure 1s22s22p63s1. When sodium atoms come together, the electron

in the 3s atomic orbital of one sodium atom shares space with the corresponding electron on a

neighboring atom to form a molecular orbital (in much the same sort of way that a covalent bond

is formed).


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The difference, however, is that each sodium atom is being touched by eight other sodium atoms

and the sharing occurs between the central atom and the 3s orbitals on all of the eight other

atoms. And each of these eight is in turn being touched by eight sodium atoms, which in turn aretouched by eight atoms and so on and so on, until you have taken in all the atoms in that lump

of sodium.

All of the 3s orbitals on all of the atoms overlap to give a vast number of molecular orbitals

which extend over the whole piece of metal. There have to be huge numbers of molecularorbitals, of course, because any orbital can only hold two electrons.

The electrons can move freely within these molecular orbitals, and so each electron becomes

detached from its parent atom. The electrons are said to be delocalized. The metal is held

together by the strong forces of attraction between the positive nuclei and the delocalizedelectrons.

This is sometimes described as an array of positive ions in a sea of electrons.

But, remember that a metal is made up of atoms and not ions.

Each positive centre in the diagram represents all the rest of the atom apart from the outerelectron, but that electron hasnt been lost it may no longer have an attachment to a particular

atom, but its still there in the structure. Sodium metal is therefore written as Na notNa+.

16

Metallic bonding in magnesium:If you work through the same argument with magnesium, you end up with stronger bonds and so

a higher melting point.

Magnesium has the outer electronic structure 3s2. Both of these electrons become delocalized, sothe "sea" has twice the electron density as it does in sodium. The remaining "ions" also have

twice the charge (if you are going to use this particular view of the metal bond) and so there will

be more attraction between "ions" and "sea".

More realistically, each magnesium atom has 12 protons in the nucleus compared with sodium's

11. In both cases, the nucleus is screened from the delocalized electrons by the same number ofinner electrons - the 10 electrons in the 1s2 2s2 2p6 orbitals.

That means that there will be a net pull from the magnesium nucleus of 2+, but only 1+ from the

sodium nucleus.

So not only will there be a greater number of delocalized electrons in magnesium, but there will

also be a greater attraction for them from the magnesium nuclei.

Magnesium atoms also have a slightly smaller radius than sodium atoms, and so the delocalizedelectrons are closer to the nuclei. Each magnesium atom also has twelve near neighbors rather

than sodium's eight. Both of these factors increase the strength of the bond still further.

Metallic bonding in transition elements:


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Transition metals tend to have particularly high melting points and boiling points. The reason is

that they can involve the 3d electrons in the delocalization as well as the 4s. The more electrons

you can involve, the stronger the attractions tend to be.

The metallic bond in molten metals:

In a molten metal, the metallic bond is still present, although the ordered structure has beenbroken down. The metallic bond is not fully broken until the metal boils. That means that boiling

point is actually a better guide to the strength of the metallic bond than melting point is. On

melting, the bond is loosened, not broken.

Secondary bonding (dispersion bonding, dipole bonding and hydrogen

bonding): Before the discussion of this type of bonding, it is necessary to have brief about theVan der Waals forces or Dispersion forces.

Inphysical chemistry, the van der Waals force (orvan der Waals interaction), named afterDutchscientistJohannes Diderik van der Waals, is the sum of the attractive or repulsive forces

between molecules (or between parts of the same molecule) other than those due tocovalent

bonds, the hydrogen bonds, or the electrostatic interactionofionswith one another or with

neutral molecules or charged molecule.

The term includes: force between two permanent dipoles (Keesom force)

force between a permanentdipole and a corresponding induced dipole (Debye force)

force between two instantaneously induced dipoles (London dispersion force).

It is also sometimes used loosely as a synonym for the totality of intermolecular forces. Van der

Waals forces are relatively weak compared to covalent bonds, but play a fundamental role infields as diverse as supramolecular chemistry, structural biology,polymer science,

nanotechnology, surface science, and condensed matter physics.

Van der Waals forces define many properties oforganic compounds, including their solubility in

polar and non-polarmedia.In low molecular weightalcohols, the hydrogen-bonding properties ofthe polarhydroxyl group dominate the weaker van der Waals interactions. In higher molecular

17

weight alcohols, the properties of the nonpolar hydrocarbon chain(s) dominate and define the

solubility. Van der Waals forces quickly vanish at longer distances between interacting

molecules.[In 2012, the first direct measurements of the strength of the van der Waals force for a single organic molecule bound to a metal surface was made

via atomic force microscopyand corroborated withdensity functional calculations.]

Secondary bonds are weak in comparison to primary bonds.

They are found in most materials, but their effects are often overshadowed by the strength of the

primary bonding.

Secondary bonds-1(dipole bonding):

Secondary bonds are not bonds with a valence electron being shared or donated. They are usually

formed when an uneven charge distribution occurs, creating what is known as a dipole (the totalcharge is zero, but there is slightly more positive or negative charge on one end of the atom than

on the other).

These dipoles can be produced by a random fluctuation of the electrons around what is normallyan electrically symmetric field in the atom.
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Once a random dipole is formed in one atom, an induced dipole is formed in the adjacent atom.

This is the type of bonding present in N2 molecules, and is known as van der Waals Bonding.Secondary bonding may also exist when there is a permanent dipole in a molecule due to an

asymmetrical arrangement of positive and negative regions.

Molecules with a permanent dipole can either induce a dipole in adjacent electrically symmetricmolecules and thus form a weak bond or they can form bonds with other permanent dipole

molecules.

Secondary bonds-2(hydrogen bonding):Hydrogen bonding is the strongest form of secondary bonding and is formed from the polar

nature of molecules containing hydrogen. The hydrogen side of the molecule is more positivethan the atom it is bonded to, allowing an attraction toform with the negative end of another molecule.

In hydrogen bonding, the hydrogen is attached directly to an element which is electronegative.

This causes the hydrogen to acquire a significant positive charge.18

The melting and boiling points of materials containing hydrogen bonding, are abnormally high,

for their atomic weights, due to hydrogen bonding.


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Notice that in each of these molecules NH3, H2O and HF above:

The hydrogen is attached directly to one of the most electronegative elements, causingthe hydrogen to acquire a significant amount of positive charge.

Each of the elements to which the hydrogen is attached is not only significantly negative,

but also has at least one "active" lone pair.

Lone pairs at the 2-level have the electrons contained in a relatively small volume of space which

therefore has a high density of negative charge. Lone pairs at higher levels are more diffuse andnot so attractive to positive things.

Variation in bonding character and Physical Properties (within a Group):The physical properties (notably, melting and boiling points) of the elements in a given group

vary as you move down the table.

The physical properties of elements depend in part on their valence electron

configurations. As this configuration remains the same within a group, physical

properties tend to remain somewhat consistent.

The most notable within-group changes in physical properties occur in groups 13, 14, and15, where the elements at the top are non-metallic, while the elements at the bottom are

metals.

19

The trends in boiling and melting points vary from group to group, based on the type of

non-bonding interactions holding the atoms together.

physical property:- A physical property is any property that is measurable whose value

describes a physical system's state.

malleable:- Able to be hammered into thin sheets; capable of being extended or shaped by

beating with a hammer or by the pressure of rollers. ductile:- Capable of being pulled or stretched into thin wire by mechanical force without

breaking.

Boiling and Melting Points:

It should be noted that some elements exist in different forms. For example, pure carbon canexist as diamond, which has a very high melting point, or as graphite, whose melting point is still

high but much lower than that of diamond.


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Different groups exhibit different trends in boiling and melting points. For groups 1 and 2, the

boiling and melting points decrease as you move down the group. For the transition metals,

boiling and melting points mostly increase as you move down the group, but they increase forthe zinc family. In the main group elements, the boron and carbon families (groups 13 and 14)

decrease in their boiling and melting points as you move down the group, whereas the nitrogen,

oxygen, and fluorine families (groups 15, 16, and 17) tend to increase in both. The noble gases

(group 18) decrease in their boiling and melting points down the group.(See page:5)For metallic species, the metallic bonding interaction (electron-sharing) becomes more difficult

as the elements get larger (toward the bottom of the table), causing the forces holding them

together to become weaker. As you move right along the table, however, polarizability and vander Waals interactions predominate, and as larger atoms are more polarizable, they tend to

exhibit stronger intermolecular forces and therefore higher melting and boiling points.

A hydrogen bondis effectively a strong example of an interaction between two permanent

dipoles. The large difference in electro-negativities betweenhydrogen and any offluorine,

nitrogen and oxygen, coupled with their lone pairs of electrons cause strong electrostatic forcesbetween molecules. Hydrogen bonds are responsible for the high boiling points.

There are four basic types of bonds that can be formed between two or more (otherwise non-

associated) molecules, ions or atoms.Intermolecular forces cause molecules to be attracted or

repulsed by each other. Often, these define some of the physical characteristics such as themelting point of a substance.

Metallic Character:

Metallic elements are shiny, usually gray or silver in color, and conductive of heat and electricity.

They are malleable (can be hammered into thin sheets) and ductile (can be stretched into wires).Some metals, such as sodium, are soft and can be cut with a knife. Others, such as iron, are very

hard. Non-metallic atoms are dull and are poor conductors. They are brittle when solid, and

many are gases at STP (standard temperature and pressure). Metals give away their valence

electrons when bonding, whereas non-metals tend to take electrons.Metallic character increases from right to left and from top to bottom. Non-metallic character

follows the opposite pattern. This is because of the other trends: ionization energy, electron

affinity, and electronegativity.

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http://en.wikipedia.org/wiki/Hydrogen_bondhttp://en.wikipedia.org/wiki/Hydrogenhttp://en.wikipedia.org/wiki/Hydrogenhttp://en.wikipedia.org/wiki/Fluorinehttp://en.wikipedia.org/wiki/Fluorinehttp://en.wikipedia.org/wiki/Nitrogenhttp://en.wikipedia.org/wiki/Oxygenhttp://en.wikipedia.org/wiki/Intermolecular_forcehttp://en.wikipedia.org/wiki/Intermolecular_forcehttp://en.wikipedia.org/wiki/Melting_pointhttp://en.wikipedia.org/wiki/Hydrogen_bondhttp://en.wikipedia.org/wiki/Hydrogenhttp://en.wikipedia.org/wiki/Fluorinehttp://en.wikipedia.org/wiki/Nitrogenhttp://en.wikipedia.org/wiki/Oxygenhttp://en.wikipedia.org/wiki/Intermolecular_forcehttp://en.wikipedia.org/wiki/Melting_point

mst,module 1, notes

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