IIT KANPUR. ACADEMICS. STUDY MATERIAL · ABOUT ME · CONTACT US · QUERY · GALLERY · ESC · PHY · PHY · MTH · MTH CHM . CHM LECTURE NOTE. COURSE TITLE: INTRODUCTORY PHYSICAL CHEMISTRY. CREDIT UNIT: PART TITLE: CHEMICAL EQUILIBRIUM. COURSE. Dr. Ron Briggs, H 2. CHM ○Introductory course - no prerequisites. ○ Serves as a ○This and future chapter notes are available as pdf files on my.

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View Notes - CHM pdf from MTH at National Open University of Nigeria. NATIONAL OPEN UNIVERSITY OF NIGERIA SCHOOL OF SCIENCE AND. Chm Introductory Inorganic Chemistry - Download as PDF File .pdf), Text File .txt) or read online. Introductory College Chemistry Course. General Chemistry Chem. pdf - Download as PDF File .pdf), Text File .txt) or read online.

A brief background on physical aspects of photochemistry will be given. Exploring and understanding of reactions that are triggered by light. Importance of light in life will be highlighted. Principles of Bonding and Reactivity in Inorganic Chemistry. Bonding principles necessary to understand the structure, stability, and fundamental reactivity of main group and transition metal inorganic compounds.

Modern Quantum Chemistry. Many-electron wave functions and operators. Hartee-Fock approximation, density functional theory, configuration interaction, and many-body perturbation theory.

Modern electrochemical techniques including voltammetry, chronocoulometry, rotating disk electrode, and ultramicroelectrodes.

Principles of Spectroscopic Techniques. Spectroscopic techniques: nuclear magnetic resonance NMR , mass spectra MS , ultraviolet UV , visible infrared IR , fluorescence, and other specialized spectroscopic techniques. Theory of nuclear magnetic resonance; Bloch equations; relaxation theory; time- domain versus frequency domain spectroscopies, and principles of multidimensional spectroscopy. Advanced Analytical Chemistry. Provides a strong foundation in the most important concepts in advanced analytical chemistry, including electrochemistry, chemical separations, and bioanalytical chemistry, and in the different classes of instrumental analytical techniques available to current chemists.

The photophysical properties of organic compounds that illustrates the fundamental principles of fluorescence. It also explains how fluorescence spectra and images can be recorded and how these powerful analytical techniques can be used to address significant problems in biology and medicine. Molecular and Supramolecular Photochemistry. A brief background on physical aspects of photochemistry will be given.

Exploring and understanding of reactions that are triggered by light. Importance of light in life will be highlighted.

Principles of Bonding and Reactivity in Inorganic Chemistry. Bonding principles necessary to understand the structure, stability, and fundamental reactivity of main group and transition metal inorganic compounds.

Modern Quantum Chemistry. Many-electron wave functions and operators. Hartee-Fock approximation, density functional theory, configuration interaction, and many-body perturbation theory. Modern electrochemical techniques including voltammetry, chronocoulometry, rotating disk electrode, and ultramicroelectrodes.

Principles of Spectroscopic Techniques. Spectroscopic techniques: nuclear magnetic resonance NMR , mass spectra MS , ultraviolet UV , visible infrared IR , fluorescence, and other specialized spectroscopic techniques. The higher the effective nuclear charge. These factors are: When an electron adds on to any atom. On the other hand the noble gases with closed sell ns2np6 configuration. The smaller the size of the atom.

Both these factors favour an increase in the force of attraction exerted by the nucleus on the extra Therefore the higher will be the electron affinity of the atom thus. Trends across periods: On moving from left to right in a period. Electronic configuration of the atom also plays an important role in determining the magnitude and sign of electron affinity.

Halogens can achieve a stable noble gas configuration by accepting just one electron. In this sub section we will learn how the electron affinity varies in the provided descriptions. This is apparently an indirect result of the small size of he atoms of these elements that is B. This is evident from the values given listed in table 3. Trend across groups: We know from the previous section that on moving down the group of s.

As explained earlier. Thus electron affinities of alkaline metals have small negative values indicating their reluctance to form an anion. Therefore even though an electron added to an atom of an element of period 2 is closer to the nucleus than one added to an atom of an element of period 3. Cl of period 3. F are however against the general trend. Answers 6. Self Assessment Question 1 2 List three factors that affect electron affinity? Why does the electrons affinity generally deceased down a group.

Explain the difference between electron affinity and eletronegativity. The electronegativity of an element is a measure of the power of an atom in a molecule to attract shared electrons to it self. In the process of formation of a bond between A and B.

Unlike ionization and electron affinity. Define the concept of electronegativity Calculate. These quantities are a measure of the tendency of isolated atoms to lose or gain electrons.

Discuss with at least eighty percent accuracy periodicity in electronegativity. In a heteronuclear diatomic molecule of the AB type however. Mulliken-Taffe and AlfredRochow electronegativity scales. As the positively develops on A. B-B the electron pair is equally shared between the atoms bonded together. Meanwhile a similar process also takes place on atom B.

He defined it on the basis of the patterns desirable in the single bond energies of elements which were derived from the thermochemical data. They continue doing so until the tendencies of both the atoms in the bonded state to attract the electron pair towards themselves balance. Partial charges will thus be generated on A and B. This formulae only gives the difference in the electronegativities of the two elements and not the absolute value assigned to a particular element. He realized that bond energy.

If one of the atoms say B. The largest electronegativity difference is that between fluorine the most electronegative element and calcium.

Knowing bond energies. Pauling assigned abitrarily a whole number value 4.

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This tendency of attracting the shared electron pair toward itsself by an atom a molecule has been termed the electronegativity of the element The concept of electonegativity was first developed by Pauling.

The trends in the variation of electronegativity are however the same. So it will be difficult to remove its electrons. Thus electronegativity XA of atom A is given by the following relationship. It is very difficult to measure electron affinity for all elements as we learned in unit 6.

The electronegativity values on Mulliken Scale are about 2. The force of attraction is expressed according to coulomb's law as: So it looses an electrons readily and has little tendency to pick up electron. Both quantities are given positive values if loss of electron involves absorption of energy and negative values if gain of electron involves release of energy. It also has a very high electron affinity. Therefore this method is not universally applicable. Thus hybrid orbitals having greater scharacter possess higher electronegativity.

He cicw W0. In turn.. All the electronegativity scales give only average values of electronegativities of element in different bonding environments. An atom which has acquired a positive charge will tend to attract electrons to it more readily than will a neutral atom. C2H4 and C2H2 has different values of electronegativity. Rochow Sole 1. An atom in sp hybridized state will be more electronegative than the same atom in Sp 2 hybridized state which will in turn be more electronegative than the same atom in SP3 hybridized state.

Re Os Ir 1. Thus the carbon atom in CH 4. So you can say that electronegativity is not a constant quantity. Hybridization also affects electronegativity because of lower energy and hence. The ability of an atom to attract electrons depends upon the charge on the atom and the hybridization of the atom.

On moving down a group of representative elements. These values are nevertheless useful in making quantitative comparisons between elements. Se and Br have greater electronegativities than would be expected by extrapolation from values for the first two elements in the respective groups.

This is as a result of the sharp increase in effective nuclear charge of these elements example from lithium to fluorine. As expected. The decrease is relatively small except between the first two elements. Electronegativities can be used to predict the value of the bonding that a compound will have. On the contrary. This is because the additional electron is being added to an inner shell which provides relatively good shielding for the outer electron from the nucleus.

The larger the difference between the electronegativities of the two elements. This is due to the insertion of transition elements because of which the effective nuclear charge of these elements is greater than that.

CHM101: Chemistry and Global Awareness (Gordon)

The much greater electronegativity of lithium row elements correlates well with their small size. The trend is similar to that of ionization energies. Mulliken or Alfred Rochow That the measurements vary according to the method used in measuring it. That the values of electronegativity vary across periods and groups of the periodic table Self Assessment Questions 1 Name the least and most electronegativity elements in the periodic table Answer 1 Caesium and francium are the least electronegative element whereas fluorine is the most electronegative element in the periodic table 6.

In this unit on the other hand. It is also the principal element in the solar atmosphere. This electronic configuration is similar to the outer electronic configuration of the alkali metals ns I. Inspite of that. The hydrogen atom consist of only one proton and one electron. However although hydrogen is very much in abundance 0. Like alkali metals. In addition. You will also be studying its position in the periodic table. I am sure that you noticed the fact that the very first element in the periodic table is hydrogen.

I am sure seeing how important hydrogen is and its peculiarity of being the first element in the periodic table. We have seen in unit 5 that the alkali metals have tight tendency of losing In this unit you will be studying some important aspects of the chemistry of hydrogen. Justify the position of hydrogen in the periodic table Describe isotopes of hydrogen Differentiate between Ortho and Para forms of hydrogen. Hydrogen is the first element of the periodic table.

Hydrogen has three different Isotopes having mass numbers 1. By picking up an electron. From the previous discussion.

Their first ionization energies are fluorine kg mol -1 chlorine kg mol -1 bromine kg mol -1 and iodine kg mol These Isotopes differ from one another in respect of the presence of neutrons. Due to its high ionization energy. Ordinary hydrogen has no neutrons. So hydrogen can be placed with either of them in the period table. Hydrogen like halogens. However conventionally. The physical properties of hydrogen. These Isotopes have the same electronic configuration and therefore their chemical properties are almost the same.

The only difference is in the rate of reactions. Some of the important physical properties of hydrogen. Deuterium oxide is used as a modei. H20 which is the oxide of hydrogen. Physical properties of H2O and D20 and also differ from each other as in the case of H2 D2 these are given in Table 3.

Hydrogen is liberated more quickly than deuterium at the cathode and the residual liquid continuously gets richer in deuterium content on prolonged electrolysis of water. Like water. It can be obtained from ordinary water which contains 0. This can be done either by fractional distillation or by electrolysis. The radiation that tritium gives off within an organism. See fig 3. When two hydrogen atoms combine to form a molecule. Tritium was first obtained synthetically by the bombardment of deuterium compounds such as ND 4 SO4 with fast deuterons.

The concentration of tritium increased by over a hundred fold when thermonuclear weapon testing began in but is now subsiding again as a result of the ban on atmospheric weapon testing. Its half life period is These different forms arise as a result of differences in the direction of nuclear spin. Para Para hydrogen has a lower internal energy than that of Ortho hydrogen.

Ortho-Para equillibria for H2. Ortho and Para forms of hydrogen Due to spin Isomerism. Hydrogen gas is an equilibrium mixture of Ortho and Para hydrogen. D2 and T2. The ration of Ortho to Para hydrogen varies with temperature as shown in fig 3. D and T two other forms.

That these two forms arise as a result of differences in spins of the two molecules that make up the H2 molecule. That hydrogen with an electronic configuration of Is' occupies a unique position in the periodic table That hydrogen has Isotopes known as deuterium tritium and the normal hydrogen we know. Hydrogen also exists in different forms thus explaining some of its properties. Deuterium and tritium also exhibit spin Isomerism and exist in.

Ni Pt. Self Assessment Question 1. That hydrogen has in addition to H. Ortho and Para forms. It is the first element in the periodic table and also exhibits properties of the Alkali metals as well as that of the hydrogen. Write T for true and 'F' for false in the given books for the following statements about Ortho and Para forms of hydrogen: However the ratio of Ortho to Para forms in deuterium and tritium is different from that in hydrogen.

Ortho and Para Hydrogen. NO2 etc. Why does hydrogen resemble alkali metals? Why are the chemical properties of Isotopes similar?

Explain the formation of hydride ion c. General and Inorganic Chemistry by J. Wilson and A. Water can be reduced to hydrogen either chemically or electrically. List and explain at least five properties of hydrogen List and explain with eighty percent degree of accuracy at least four uses of hydrogen.

In this unit we shall be studying the ways hydrogen is manufactured. Water is a natural abundant source for the manufacture of hydrogen.

List at least 3 methods used in the manufacture of hydrogen Explain using appropriate equations and examples what takes placeduring the process of the manufacture of hydrogen. These reactions. At higher temperatures. One example of simple cracking reaction is cracking of propane. In both cases above CO is converted to CO 2. Electrolysis of acidified water using platinum electrodes is a convenient source of hydrogen and oxygen.

On a large scale. Hydrogen obtained by electrolysis of water is relatively expensive because of the cost of electrical energy. It is colourless. The atomic hydrogen produced. During electolysis. Most of the transition metals catalyse the combination reaction of hydrogen. The reaction can be explosively violent with alkali metals e. H2 does dissociate.

The hydrogen molecule is ii Thermally stable and has little tendency to dissociate at normal temperatures. Atomic hydrogen is produced by passing ordinary hydrogen through electric arc maintained between two electrodes.

It reacts with almost all elements except the noble gases. H2O and HF. The atoms set free are carried away by a stream of incoming hydrogen gas. CaH 2 xi With non-metals it forms covalent hydrides. This principle is utilized in the making of the 'atomic hydrogen welding torch' see fig 3. These free atoms recombine at once on coming in contact with a metallic surface liberating a large amount of heat and thus arising temperature of the metal to say — k.

Such metals include Mo and W. In the presence of catalysts such as finely divided nickel. Atomic hydrogen welding torch xiii Hydrogen is easily oxidized to water and. The one in fig 3. In a fuel cell. This is sometimes called "cold combustion" A hydrogen oxygen fuel cell may be have an alkaline of acidic electrolyte. The largest single use of hydrogen is in the syntheses of ammonia which is used in the manufacture of nitric acid and nitrogenous fertilizers.

Hydrogen is used in the hydrogenation of vegetable of oils and the manufacture of methanol In space crafts. This is because we are using the same reactants at the electrodes in both cases. A hydrogen-oxygen fuel cell with KOH electrolyte and porous carbon electrodes The half cell reactions are given below: Firstly in a fuel cell unlike in the dry cell or storage battery which requires recharging also.

Fuel cells have several advantages over other sources of energy. Nuclear and hydro electric power cannot meet the world's energy needs. Hydrogen as a fuel has many advantages over the conventional fossil fuels and electric power. Hydrogen however has the following disadvantages viz: Hydrogen like electricity is a secondary source of energy because it is produced using energy from a primary source such as coal.

Preparation of hydrogen through electrolysis is not economical at present in fact more energy has to be spent in electrolysis of water than what can be liberated by burning hydrogen as a fuel. In addition to solar power. It releases greater energy per limit weight of fuel in comparison to gasoline and other fuels. Unlike electricity hydrogen can be stored and used when needed. It is now clear that world reserves of fossil fuels like coal.

Hydrogen can be transported as a gas in high pressure pipelines.


Decomposition of water by solar energy in presences of catalysis is known as photochemical decomposition of water. If this process can be made industrial. With these facts in mind there is now an active search for alternative source of energy. It is available in unlimited quantities in sea water. It is pollution free because the major product of its combustion is water with only traces of nitrogen oxides. Using catalysts scientists in France have been able to achieve the efficient decomposition of water under visible and ultra violet light.

Complete the following chemical reactions which take place during the formation of hydrogen. They are: The single most important use of hydrogen is in the manufacture of ammonia which is used in the production of nitric acid and nitrogenous fertilizers.

That hydrogen is made use of in the production of fuel ceils for space crafts. Exercises That the largest use of hydrogen is in the manufacture of ammonia.

Since water is abundant. Hydrogen is also used in space crafts as a source of fuel cell. That hydrogen is also used in extraction of metal and hydrogenation of vegetable oils.

That the major source for the manufacture of hydrogen is water That there are two primary methods for the manufacture of hydrogen a. This is non economical because of the cost of electricity. Formation of methanol from coal i ii Reduction of methyl cyanide iii Conversion of Oleic acid into stearic and iv Reduction of ammonium molybdate to molybdenum 7. Ionic or salt like or salric hydrides Covalent or molecule hydrides Metallic or non-stoichiometric hydrides 3.

In this unit you will be studying the formation of hydrides. As electronegativity of the element increases the stability of the hydrides also increases. Hydrogen combines with a number of elements to form hydrides. As already pointed out at the beginning of this unit. These hydrides are formed by transfer of an electron from the metal to the hydrogen atom.

Chemistry (CHM)

Hydride ion is a peculiar chemical species and in contrast to proton which has small size. We also saw in unit 9. You also learned that it forms ionic and covalent hydrides with metals and non metals respectively. Three types of hydride compounds are formed depending upon the electronegativity of the element. There are classified into: It is larger than any of the negative ions except iodide. List the three classes of hydrides Discuss for each of the different types of hydrides and their properties.

They are powerful reducing agents especially at high temperatures e. They have high melting points and conduct electricity in liquid state.

The complex hydrides are frequently used in the reduction of aldehydes. Alkali and Alkaline-earth metal of groups 1 and 2 are sufficiently electropositive and force the hydrogen atom to accept an electron to form the hydride ion. The bonds formed in this class of hydrides are mostly covalent in character but in some cases. Ionic hydrides are formed by heating metals in hydrogen at k. Their density is higher than that of the metal.

Ionic hydrides are white crystalline solids.


The covalent hydrides can be prepared either by direct reaction of nonmetals with hydrogen under suitable conditions or by the reaction of H2 O or acids or nitrides, carbides, bonides, silicides, stanides of alkali and alkaline earth metals or by the reduction of halides.

These are illustrated by the following reactions. These hydrides have molecular lattice made up of individual saturated covalent molecules, with only weak Vander Waals forces and in some cases along with hydrogen bonds.

This accounts for their softness, low melting and boiling points, their volatility and lack of conductivity. Some covalent hydrides are unstable in the presence of air, e.

Some covalent hydride hydrides of groups 2 and 13 are electron deficient. These have structures between ionic and covalent hydrides. These are either dimeric, e. Most of these have metallic appearance and are less dense than the parent metal. They all conduct heat and electricity though not as well as the parent metal. They are almost always non-stoichiometric, being deficient in hydrogen.

For example, Ti H5,8 VH 0. Most of these hydrides are stable to water up to K but are quantitatively decomposed by acids and show some reducing properties. Formerly these hydrides were formed as interstitial compounds in which hydrogen was through to be accommodated in the interstices in the metal lattice producing distortion but no change in its type. But recent studies have shown that except for hydrides of nickel palladium, cerium and actinium, other hydrides of this class have lattice of a type different from that of the parent metal.

For example, the hexagonal close packed lattice of some lanthanides is transformed to a facecentred cubic lattice in their dihydrides.

As pointed out earlier, these interstitial hydrides are poorer conductors of electricity, exhibit less Para magnetism and are more brittle than the parent metal. These characteristics suggest that hydrogen is present in the metal lattice as hydrogen atoms rather than as hydrogen molecules.

The single electron of hydrogen is paired with an electron of the metal, thereby reducing the extent of metallic bonding.

Breaking of the H-H bond is in agreement with the fact that there metals catalyse reactions of hydrogen. That hydrogen leads with certain metals to form hydrides That these hydrides are classified into three viz ionic or salt — like hydrides, covalent hydrides and metallic hydrides That ionic hydrides are powerful reducing agents, covalent hydrides are soft have low melting points and are poor conductors of electricity.

Exercises Silt is an example of which of the following type of hydrides 1. Answer d Covalent hydride List 3 properties of Ionic hydrides Answer: They have high melting points i ii They conduct electricity in fused state liberating hydrogen at the anode iii Their density is higher than that of the metal. This bond is known as hydrogen bond. Define hydrogen bonding List and discuss the two types of hydrogen bonding List and discuss the effect of bonding on boiling and melting points Discuss the effect of bonding on the solubility of a substance Briefly discuss.

When hydrogen is covalently bonded to a highly electronegative element like F. When this happens the hydrogen. Intermolecular Hydrogen Bonding In this case. H-0 and H—F bonds. There are two types of hydrogen bonding.

Some common examples of intermolecular hydrogen bonding occurring between the molecules of the same compound are HF. The hydrogen bond energy is only about 7 — 59 kg mol -1 compared to the normal covalent bond energy of — kg mol l for H--N.

H 2 O alcohols etc. These are: Salkylaidohyds aNItrophod You have seen above. As a consequence of this. This anomaly is explained on the basis of hydrogen bond formation. In this section we shall discuss its effect on the melting point. HF in Groupl7 and water have abnormally high melting and boiling points as compared to other hydrides in their respective groups in the periodic table.

NH 3 in Group It does however show the usual intermolecular hydrogen bonding. In compounds where the molecules are linked by hydrogen bonds. Bolling point Te mpera ture K 50 0 3 2 5 4 Period of combining element 3 4 Fig. Due to the intramolecular chelated structure. We can also say that the polarizing power of a cation is proportional to the ratio of its charge to its size.

It is mainly because of this reason that many covalent hydrides H-X are acidic in aqueous solution. This ratio is known as the ionic potential of the cation.

As the hydrogen cation. That hydrogen bond is defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom of another molecule. Variations in boiling and melting points and solubility of substances are explainable by the existence of hydrogen bonds. As result of this polarizing power for protons. H Enthalpy of formation of these aquated proton species is large kg moi l.

They are generally found associated with other molecules. That there are two types of hydrogen bonds viz: Intermolecular hydrogen bond i ii Intra molecular hydrogen bond That the hydrogen bonding in substances affect their melting points.

HF and NH 3? Why is H2 O a liquid and H 2 S a gas at room temperature? Why are group 2 metals harder and have higher melting points as compared to the Group 1 metals? Since Group 2 elements have two valence electrons. What is the effect of hydrogen bonding on he properties of H Thus they are harder and have higher meeting points as compared to the Group 1 elements.

Exercises 1. Which distance is shorter? In earlier units. Group 1 elements consists of Li. You also recall how in the last few units we saw that the properties of the elements in a periodic table are indeed a periodic function of their atomic numbers as stated by the periodic law. We also saw that elements belonging to the same groups have the same properties.

The elements of group 1 and 2 are called the S-block elements because the outermost electron s in these elements occupy the S-orbital. You will also be studying their physical properties. You recall that the end product of that effort is the arrangement of the elements according to their atomic numbers. In this unit you will be studying the elements of group 1.

C and Fr. Potassium in carnallite. Potassium is obtained by the reduction of its chloride with sodium vapor. Francium being a radio active element with a very short half life period Rubidium and Caesium. They occur in the combined form in the earth's crust in the following relative abundance: Sodium 2. Lithium occurs in alumino silicate rocks. Na 3A1F6. OH 2 Sodium in rock salt. Rubidium and calcium are rare elements and generally occur in small quantities along with other alkali metals.

Sodium as sodium chloride is the most abundant metal in sea water M1. This reduction by Na appears to be contrary to the normal order of reactivity.

Caesium has the distinction of being the metal from which electrons are ejected most easily on exposure to light. Production of francium as a result of a decay. Sodium in polyethylene enclosed cables is used in some underground high voltage transmission applications.

This phenomenon is called photoelectric effect. These are sodium vapour lamps and the light from them can penetrate fog well. Rubidium and caesium salts are obtained during the recrystallisation of other naturally occurring alkali metal salts.

Because of the high specific heat and thermal conductivity. However due to their highly reactive nature they cannot be used for this purpose.The small difference in the conductivity is due to the value of the metal itself. Mendeleev — Ordered elements by atomic mass.

It also explains how fluorescence spectra and images can be recorded and how these powerful analytical techniques can be used to address significant problems in biology and medicine. Ionic hydrides are white crystalline solids.

The electronic configurations of isolated atoms of elements are usually verified experimentally by a detailed analysis of atomic spectra. Whenever you use an instrument to compare a quality of an object to a standard. These are the atomic radii decreases along a period and generally increase down a group in the long form of the periodic table see fig 3. At high temperatures. Thus the ionization energies of the alkali metals are the lowest and those of the noble gases are the highest in their respective periods.

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