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.
