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Atomic Structure & Crystallization Presentation

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Unit – I – Atomic structure and crystallization
Dr. G. Viswanathan
Assistant Professor in Chemistry
Bannari Amman Institute of Technology
UNIT I
ENGINEERING CHEMISTRY – I
18ME103
ATOMIC STRUCTURE AND CRYSTALLIZATION
Atomic structure: Introduction to fundamental concepts - dual nature of the electrons periodic table - types of atomic interaction (ionic, covalent, coordinate covalent, metallic
and Vanderwaals interactions). Metallic crystal structure - ceramic crystal structure polymer crystal structure.
UNIT II
PHASE RULE
Phase - component - degree of freedom - solubility limit - Gibbs phase rule - phase
diagram - phase equilibrium applications - one component system (water system).
Reduced phase rule: Two component systems (lead and silver system and Fe-Fe3C
diagrams).
UNIT III
FERROUS AND NON-FERROUS ALLOYS
Alloys: Purpose of alloying - function and effects of alloying elements - properties of alloys
- classification of alloys. Composition, types, properties and applications of ferrous alloys
(Steel, cast iron and stainless steel), Non-ferrous alloys (Aluminium, copper, magnesium
and nickel).
Continue…
UNIT IV
HEAT TREATMENT
Heat treatment of steel: Annealing - stress relief - recrystallization and spheroidizing normalizing - hardening - tempering of steel - isothermal transformation diagram (TTT
diagram) - cooling curves - carburizing - nitriding - cyaniding - carbonitriding - flame and
induction hardening.
UNIT V
SPECTROSCOPY
Beer-Lamberts law - Electromagnetic spectrum - electronic - vibrational - rotational
transitions. Principle - instrumentation (Block diagram) - applications of UV visible - IR
spectroscopy. Spectrophotometric estimation of iron (colorimetry).
Unit – I – Atomic structure and crystallization
Topic: Introduction to fundamental concepts, dual nature of
the electrons and periodic properties
General Objective
Students will be able to understand the fundamental concepts of atomic
structure and apply the periodicity of elements to define their
properties.
Specific Objectives (SO):
Students will be able to
•SO1: Illustrate the progress of atomic models and quantum
number for the explanation of atomic structure (U, C)
•SO2: Justify the duel nature (wave and particle) of electrons
and outline the de Broglie’s concept (An, C).
•SO3: Compare the periodic properties such as atomic radii,
ionic radii, electron affinity, electro negativity and
ionization energy of various groups of elements in
periodic table (An, C).
1. Count the total number of materials in the picture
2. How many of the things used as packaging the materials
The most abundant element in the earth’s crust
is oxygen.
The Indian Sage who developed
Atomic Theory 2,600 years ago
Acharya Kanad was born
in 600 BC in Prabhas
Kshetra (near Dwaraka) in
Gujarat, India. His real
name was Kashyap.
People began calling him ‘Kanad’, as ‘Kan’ in Sanskrit means ‘the smallest particle’
Democritus about atom
• Atoms are small hard particles.
• Made of a single material that’s formed
into different shapes and sizes.
• They are always moving
• They form different materials by joining
together.
Democritus was an ancient Greek philosopher who lived
from 460 - 370 B.C.
John Dalton 1776-1844
• Dalton a chemist and school teacher brings back
idea of the atom.
• He performed many experiments to study how
elements join together to form new substances.
• He found that they combine in specific ratios
(remember the electrolysis of water) and he
supposed it was because the elements are made
of atoms.
John Dalton theory of atom
1. Elements are composed of small indivisible particles
called atoms.
2. Atoms of the same element are identical. Atoms of
different elements are different.
3. Atoms of different elements combine together in simple
whole number ratios to create compounds.
4. In a chemical reaction, atoms are rearranged, but not
changed.
Was Dalton right?
Elements are
composed of small
indivisible particles
called atoms.
Subatomic particles –
electrons, protons,
neutrons, and more
Atoms of the same
element are identical.
Atoms of different
elements are different.
No, isotopes are atoms
that have the same
number of protons but
a different number of
neutrons
Was Dalton right?
Atoms of different
Yes! He was right!
elements combine
together in simple
whole number ratios to
create a compound.
In a chemical reaction, Yes! He was right!
atoms are rearranged,
but not changed.
J. J. Thomson (1903)
• Cathode Ray Tube
Experiments
– beam of negative particles
• Discovered particles smaller
than the atom!
Thomson’s Cathode Ray Experiment
Stream of electrons
are attracted to
positively charged
plate here.
The Plum Pudding Model
• Proved that the atom was divisible and that all atoms
contain electrons.
• He proposed that the atom was a sphere of positively
charged material. Spread throughout the atom were
the negatively charged electrons similar to plums in a
pudding or chocolate chips in ice cream.
Thomson did not know how
the electrons in an atom
were arranged. He believed
they
were
mixed
throughout an atom.
Ernest Rutherford (1871 – 1937)
• Awarded the Nobel Prize in Chemistry for
his discovery of alpha particles, positively
charged particles emitted from radioactive
elements.
• Was a student of J.J. Thomson but disagreed
with the “Plum Pudding Model”
• Devised an experiment to investigate the
structure of positive and negative charges in
the atom.
Rutherford’s Revised Atomic Theory (1911)
Result: Most of the positively charged particles went straight through
the gold foil.
Atomic Theory: Most of the matter of the atom is found in a very small
part of the atom. This is called the nucleus of the atom. It is very
tiny and extremely dense.
Result: Some of the positively charged particles were deflected or even
bounced back.
Atomic Theory: Like charges repel so the nucleus must have a positive
charge. If electrons have a negative charge they could not be in a
positively charged nucleus. Electrons must surround the nucleus at
a distance.
Result: The diameter of the nucleus is 100,000 times smaller than the
diameter of the entire gold atom.
Atomic Theory: Atoms are mostly empty space with a tiny, massive
nucleus at the center .
Niels Bohr (1913)
• Met with J.J. Thomson but
didn’t impress him
• Worked with Rutherford
and liked his model of the
atom
• Incorporated idea of
quantum mechanics into
the Rutherford model
Erwin Schrödinger (1926)
• Treats electrons as waves
• Tells us the probability of finding an electron
at any given location at any given moment
• Electron cloud model
– Atomic orbital: region around the nucleus where
electrons are likely to be found.
Electron Cloud Model (orbital)
• dots represent probability of finding an enot actual electrons
The concept of Quantum numbers was introduced to
distinguish the orbital on the basis of their size, shape and
orientation in space by using principal, azimuthal, magnetic
and spin quantum numbers. From the study of quantum
numbers, various rules are put forward for filling of electrons
in various orbitals as follows.
Hunds rule of maximum multiplicity.
Hund's rule: every orbital in a subshell is singly
occupied with one electron before any one orbital is
doubly occupied, and all electrons in singly
occupied orbitals have the same spin.
DUAL PROPERTY OF AN ELECTRON
Experiments to prove particle and wave property of Electrons
a) Verification of Wave character
i) Davisson and Germer’s Experiment
In 1927 Davisson and Germer observed that, a beam of electrons obtained from a
heated tungsten filament is accelerated by using a high positive potential. When this
fine beam of accelerated electron is allowed to fall on a large single crystal of nickel,
the electrons are scattered from the crystal in different directions. The diffraction
pattern so obtained is similar to the diffraction pattern obtained by Bragg’s experiment
on diffraction of X-rays from a target in the same way. Since X-rays have wave
character, therefore, the electrons must also have wave character associated with them.
Moreover, the wave length of the electrons as determined by the diffraction
experiments were found to be in agreement with the values calculated from de-Broglie
equation. From the above discussion, it is clear that an electron behaves as a wave.
ii) Thomson’s experiment
G.P. Thomson in 1928 performed experiments with
thin foil of gold in place of nickel crystal. He observed
that if the beam of electrons after passing through the
thin foil of gold is received on the photographic plate
placed perpendicular to the direction of the beam, a
diffraction pattern is observed as before. This again
confirmed the wave nature of electrons.
b) Verification of the particle character
The particle character of the electron is proved by the
following different experiments:•a light paddle wheel placed in the path of cathode rays such
that the cathode rays strike the blades of upper half, it starts to
rotate. Hence, cathode rays are consists of particle.
• Experiments such as J. J. Thomson’s experiment for
determination of the ratio of charge to mass (i.e. e/m) and
Milliken oil drop experiment for determination of charge on
electron also show that electron has particle character.
•The phenomenon of Black body radiation and Photoelectric
effect also prove the particle nature of radiation.
Metals
Tend to be lustrous,
malleable, ductile,
and good conductors
of heat and electricity.
• Compounds formed
between metals and
nonmetals tend to be
ionic.
• Metal oxides tend to
be basic.
Properties of Non-Metals
• Non-metals are poor conductors of heat
and electricity.
• Non-metals are not ductile or malleable.
• Solid non-metals are brittle and break
easily.
• They are dull.
• Many non-metals are gases.
Sulfur
• Tend to gain electrons in
reactions with metals to acquire
noble gas configuration.
Properties of Metalloids
• Metalloids (metal-like) have
properties of both metals
and non-metals.
• They are solids that can be
shiny or dull.
• They conduct heat and
electricity better than nonmetals but not as well as
metals.
• They are ductile and
malleable.
Silicon
Metalloids
• Have some
characteristics of
metals, some of
nonmetals.
• For instance, silicon
looks shiny, but is
brittle and fairly poor
conductor.
Families
• Columns of elements are called
groups or families.
• Elements in each family have
similar but not identical
properties.
• For example, lithium (Li),
sodium (Na), potassium (K), and
other members of family IA are
all soft, white, shiny metals.
• All elements in a family have the
same number of valence
electrons.
Periods
• Each horizontal row of
elements is called a period.
• The elements in a period are
not alike in properties.
• In fact, the properties change
greatly across even given row.
• The first element in a period
is always an extremely active
solid. The last element in a
period, is always an inactive
gas.
Hydrogen
• The hydrogen square sits atop Family AI, but
it is not a member of that family. Hydrogen is
in a class of its own.
• It’s a gas at room temperature.
• It has one proton and one electron in its one
and only energy level.
• Hydrogen only needs 2 electrons to fill up its
valence shell.
Alkali Metals
• The alkali family is found in the
first column of the periodic
table.
• Atoms of the alkali metals have
a single electron in their
outermost level, in other
words, 1 valence electron.
• They are shiny, have the
consistency of clay, and are
easily cut with a knife.
Alkali Metals
• Alkali metals (except Li) react with oxygen to form
peroxides.
• K, Rb, and Cs also form superoxides:
K + O2  KO2
• Produce bright colors when placed in flame.
Alkali Metals
• Found only as compounds in nature.
• Have low densities and melting points.
• Also have low ionization energies.
Alkali Metals
• They are the most reactive
metals.
• They react violently with water.
• Alkali metals are never found as
free elements in nature. They
are always bonded with
another element.
• Soft, metallic solids.
• Name comes from
Arabic word for ashes.
Alkaline Earth Metals
• They are never found uncombined in nature.
• They have two valence electrons.
• Alkaline earth metals include magnesium and
calcium, among others.
Alkaline Earth Metals
• Have higher densities and melting points
than alkali metals.
• Have low ionization energies, but not as low
as alkali metals.
• Be does not react with water, Mg reacts only
with steam, but others react readily with water.
• Reactivity tends to increase as go down group.
Boron Family
• The Boron Family is named
after the first element in the
family.
• Atoms in this family have 3
valence electrons.
• This family includes a
metalloid (boron), and the
rest are metals.
• This family includes the most
abundant metal in the earth’s
crust (aluminum).
Oxygen Family
• Atoms of this family have 6
valence electrons.
• Most elements in this family
share electrons when forming
compounds.
• Oxygen is the most abundant
element in the earth’s crust. It
is extremely active and
combines with almost all
elements.
Group 6A
• Oxygen, sulfur, and selenium are nonmetals.
• Tellurium is a metalloid.
• The radioactive polonium is a metal.
Carbon Family
• Atoms of this family have 4
valence electrons.
• This family includes a nonmetal (carbon), metalloids, and
metals.
• The element carbon is called
the “basis of life.” There is an
entire branch of chemistry
devoted to carbon compounds
called organic chemistry.
Halogen Family
• The elements in this family
are fluorine, chlorine,
bromine, iodine, and astatine.
• Halogens have 7 valence
electrons, which explains why
they are the most active nonmetals. They are never found
free in nature.
Halogen atoms only need to gain 1
electron to fill their outermost energy
level.
They react with alkali metals to form salts.
Group VIIA: Halogens
• Prototypical nonmetals
• Name comes from the Greek halos and gennao:
“salt formers”
Group VIIA: Halogens
• Large, negative electron
affinities
– Therefore, tend to oxidize
other elements easily
• React directly with metals to
form metal halides
• Chlorine added to water
supplies to serve as
disinfectant
Noble Gases
•
•
•
•
•
Noble Gases are colorless gases that are extremely un-reactive.
One important property of the noble gases is their inactivity. They are
inactive because their outermost energy level is full.
Because they do not readily combine with other elements to form
compounds, the noble gases are called inert.
The family of noble gases includes helium, neon, argon, krypton, xenon,
and radon.
All the noble gases are found in small amounts in the earth's
atmosphere.
Group VIIIA: Noble Gases
• Astronomical ionization energies
• Positive electron affinities
– Therefore, relatively unreactive
• Monatomic gases
Group VIIIA: Noble Gases
• Xe forms three
compounds:
– XeF2
– XeF4 (at right)
– XeF6
• Kr forms only one stable
compound:
– KrF2
• The unstable HArF was
synthesized in 2000.
Transition Metals
• Transition Elements
include those elements in
the B families.
• These are the metals you
are probably most
familiar: copper, tin, zinc,
iron, nickel, gold, and
silver.
• They are good conductors
of heat and electricity.
Transition Elements
• Transition elements have properties similar to
one another and to other metals, but their
properties do not fit in with those of any other
family.
• Many transition metals combine chemically
with oxygen to form compounds called oxides.
Rare Earth Elements
• The thirty rare earth
elements are composed of
the lanthanide and actinide
series.
• One element of the
lanthanide series and most
of the elements in the
actinide series are called
trans-uranium, which means
synthetic or man-made.
Periodic Properties
of the Elements
Sizes of Atoms
The bonding atomic
radius is defined as
one-half of the
distance between
covalently bonded
nuclei.
Atomic radius decreases from left to right within a period. This is
caused by the increase in the number of protons and electrons across
a period. One proton has a greater effect than one electron; thus,
electrons are pulled towards the nucleus, resulting in a smaller radius.
Atomic radius increases from top to bottom within a group. This is
caused by electron shielding.
Atomic radius
Atomic size gradually decreases from left to right across a period of elements. This is
because, within a period or family of elements, all electrons are added to the same
shell. However, at the same time, protons are being added to the nucleus, making it
more positively charged. The effect of increasing proton number is greater than that
of the increasing electron number; therefore, there is a greater nuclear attraction.
This means that the nucleus attracts the electrons more strongly, pulling the atom's
shell closer to the nucleus. The valence electrons are held closer towards the nucleus
of the atom. As a result, the atomic radius decreases.
Sizes of Ions
Ionic size depends
upon:
Nuclear charge.
Number of
electrons.
Orbitals in which
electrons reside.
Sizes of Ions
• Cations are smaller
than their parent
atoms.
– The outermost
electron is removed
and repulsions are
reduced.
Sizes of Ions
• Anions are larger
than their parent
atoms.
– Electrons are added
and repulsions are
increased.
Sizes of Ions
• Ions increase in size as
you go down a
column.
– Due to increasing value
of n.
Sizes of Ions
• In an isoelectronic series, ions have the same
number of electrons.
• Ionic size decreases with an increasing nuclear
charge.
Ionization Energy
Ionization energy is the energy required to remove an
electron from a neutral atom in its gaseous phase.
The ionization energy of the elements within a period
generally increases from left to right. This is due to valence
shell stability.
The ionization energy of the elements within a group
generally decreases from top to bottom. This is due to
electron shielding. The noble gases possess very high
ionization energies because of their full valence shells as
indicated in the graph. Note that helium has the highest
ionization energy of all the elements.
Trends in First Ionization Energies
Trends in First Ionization Energies
• going down a
column, less
energy to remove
the first electron.
– For atoms in the
same group, Zeff is
essentially the
same, but the
valence electrons
are farther from
the nucleus.
Ionization Energy
• It requires more energy to remove each successive
electron.
• When all valence electrons have been removed, the
ionization energy takes a quantum leap.
Electron Affinity
electron affinity is the ability of an atom to accept an electron.
Unlike electronegativity, electron affinity is a quantitative
measurement of the energy change that occurs when an electron
is added to a neutral gas atom. The more negative the electron
affinity value, the higher an atom's affinity for electrons.
Energy change accompanying addition of
electron to gaseous atom:
Cl + e−  Cl−
Trends in Electron Affinity
Electron affinity increases from left to right within a period. This
is caused by the decrease in atomic radius.
Electron affinity decreases from top to bottom within a group.
This is caused by the increase in atomic radius.
Electronegativity Trends
Electronegativity can be understood as a chemical property describing
an atom's ability to attract and bind with electrons.
From left to right across a period of elements, electronegativity
increases. If the valence shell of an atom is less than half full, it
requires less energy to lose an electron than to gain one.
From top to bottom down a group, electronegativity
decreases. This is because atomic number increases down a group,
and thus there is an increased distance between the valence
electrons and nucleus, or a greater atomic radius.
Mind Map
Type of atomic interactions
• A bond results from the attraction of nuclei for
electrons
– All atoms trying to achieve a stable octet
• The octet rule refers to the tendency of atoms
to prefer to have eight electrons in the valence
shell.
In general, we do not consider d or f electrons. Only the s and p
electrons are involved in the octet rule, an octet in these atoms
corresponds to an electron configurations ending with s2p6
Octet Rule = atoms tend to gain, lose or share electrons so as to
have 8 electrons
C would like to Gain 4 electrons
N would like to Gain 3 electrons
O would like to Gain 2 electrons
IONIC BOND
Complete transfer of electron from one atom to another atom
makes bond formation between the two ions by the transfer
of electrons.
Examples; NaCl, CaCl2, K2O
Electron from Na is transferred to Cl, this causes a charge
imbalance in each atom. The Na becomes (Na+) and the Cl
becomes (Cl-), charged particles or ions.
Electron Dot Structures
Symbols of atoms with dots to represent the valence-shell
electrons
1
2
13
14
15
16
17
18
H
He:

Li Be


B 


C


Na Mg


Al

N



O




 Si 
 P
S





: F  :Ne :




:Cl  :Ar :


COVALENT BOND
bond formed by the
sharing of electrons
Covalent Bond
• Between nonmetallic elements of similar electronegativity.
• Formed by sharing electron pairs
• Stable non-ionizing particles, they are not conductors at any
state
• Examples; O2, CO2, C2H6, H2O, SiC
Bonds in all the polyatomic ions and diatomics are all covalent bonds
NONPOLAR
COVALENT BONDS
when electrons are
shared equally
H2 or Cl2
2Covalent
bonds- Two atoms share one or more pairs
of outer-shell electrons.
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)
POLAR COVALENT
BONDS
when electrons are
shared but shared
unequally
H2O
- water is a polar molecule because oxygen is more electronegative than hydrogen, and
therefore electrons are pulled closer to oxygen.
METALLIC BOND
bond found in
metals; holds metal
atoms together
very strongly
Metallic Bond
• Formed between atoms of metallic elements
• Electron cloud around atoms
• Good conductors at all states, lustrous, very high melting
points
• Examples; Na, Fe, Al, Au, Co
Metallic bond, A Sea of Electrons
Metals Form Alloys
Metals do not combine with metals. They form
Alloys which is a solution of a metal in a metal.
Examples are steel, brass, bronze and pewter.
Intermolecular attractions
• Attractions between molecules
– van der Waals forces
• Weak attractive forces between non-polar
molecules
– Hydrogen “bonding”
• Strong attraction between special polar
molecules
van der Waals
• Non-polar molecules can exist in liquid and
solid phases
because van der Waals forces keep the molecules
attracted to each other
• Exist between CO2, CH4, CCl4, CF4, diatomics
and monoatomics
van der Waals periodicity
• increase with molecular mass.
• increase with closer distance between
molecules
– Decreases when particles are farther away
Hydrogen “Bonding”
• Strong polar
attraction
– Like magnets
• Occurs ONLY
between H of one
molecule and N, O,
F of another
H “bond”
H is shared between
2 atoms of OXYGEN or
2 atoms of NITROGEN or
2 atoms of FLUORINE
Of
2
different
molecules
Why does H “bonding” occur?
• Nitrogen, Oxygen and Fluorine
– small atoms with strong nuclear charges
• powerful atoms
– very high electronegativities
Intermolecular forces dictate chemical
properties
• Strong intermolecular forces cause high
b.p., m.p. and slow evaporation (low vapor
pressure) of a substance.
Which substance has the highest
boiling point?
• HF
• NH3
• H2O
Fluorine has the highest e-neg,
SO
HF will experience the
strongest H bonding and 
• WHY?
needs the most energy to
weaken the i.m.f. and boil
References
1. P. C. Jain and Monica Jain, Engineering
Chemistry, 16th Edition, Dhanpat Rai Publisher,
New Delhi, 2016.
2. Sashi Chawla, Text Book of Engineering
Chemistry, Dhanpat Rai Publications, New Delhi,
2013.
3. J. C. Kuriacose and J. Rajaram, Chemistry in
Engineering & Technology, Vol. 1& 2, TMH,
2009.
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