inorganic chemistry 1

Created by

Dra.Hj. Bayharti, MSc

Miftahul Khair,S.Si, MSc

Dra. Andromeda, MSi

INORGANIC CHEMISTRY I(3 credit semester hours )

http://kimia.unp.ac.id

CHEMISTRY DEPT. FACULTY OF MATHEMATIC AND SCIENCES

STATE UNIVERSITY OF PADANG

Contents

• - Introduction to inorganic chemistry.

• - Atomic structure, development of atomic theory especially in atomic model of wave mechanics, electron configuration.

• - Periodical Table, Effective atomic Charge and relation of

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Chemistry Department

• - Periodical Table, Effective atomic Charge and relation of that with periodical properties, periodical properties of the elements.

• -Ionic compound, formation of ionic compound, lattice energy, Born Haber cycle, ionic radii and properties of ionic compounds

• - Covalent compound, properties, octet rule, resonance, formal charge, dipole moment, VBT and MOT Theory.

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• -Molecular structure, VSEPR and hybrid.

• - Coordination Compound, introduction, nomenclature, coordination number, ligands, theory of coordination ( VBT and CFT )

References

• Miessler, G.L. and Tarr, D.A., (1999), “Inorganic Chemistry”, Prentice-Hall International, Inc, London

• Manku, G.S. (1980), “Theoretical Principles of Inorganic Chemistry”, Tata-McGraw Hill Publishing Company Limited, New Delhi.Gilreath,E.S., (1963), “Fundamental Concepts of

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• Gilreath,E.S., (1963), “Fundamental Concepts of Inorganic Chemistry”, McGraw-Hill Book Company, Tokyo.

• Huheey, J. E., Keiter, E. A. and Keiter, R. L., (1993), “Inorganic Chemistry (Principles of Structure and Reactivity)”, Ed. 4., Harper Collins College Publishers

Lecture 1Introduction to Inorganic

Chemistry

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Chemistry

What is Inorganic Chemistry ?

1. If organic chemistry is defined as the chemistry of hydrocarbon compounds and their) derivatives, inorganic chemistry can be described broadly as the chemistry of "everything else." This includes

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the chemistry of "everything else." This includes all the remaining elements in the periodic table, as well as carbon, which plays a major role in many inorganic compounds. (Gary L. Miessler and Donalt A. Tarr in Inorganic Chemistry

Inorganic Chemistry is the experimental

Inorganic Chemistry is any phase of chemistry of interest to inorganic chemist (James E. Huheey in Inorganic Chemistry., principle of structure and reactivity. This mean that matter of inorganic chemistry is broad and overlap with other chemistry discipline

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Inorganic Chemistry is the experimentalinvestigation and theoretical interpretation ofthe properties and reactions of all elements andtheir compounds except the hydrocarbons andmost of their derivatives (T. Moeller

chemistry discipline

Contras with organic Chemistry

Some comparison between organic and inorganic chemistry are in order1. Number of bonds2. Kind of bonds3. Location of hydrogen and alkyl

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3. Location of hydrogen and alkyl 4. Geometry of compound 5. Number of elements

Number of bonds

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Chemistry DepartmentBoth organic and Inorganic chemistry have a single and triple bond

Kind of bond

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Organic chemistry has sigma and phi bond, inorganic chemistry has sigma, phi and delta bonds, because metal atom has d orbital

Location of hydrogen

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In organic compound, hydrogen is nearly bonded to single carbon. In organic compound, especially encountered as a bridging atom between two or more others atoms. Alkyl groups may also act as bridges in inorganic compounds

Coordination number and geometry

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Carbon is usually limited to a maximum coordination number of four (a maximum of four atoms bonded to carbon, as in CH4) and the arrangement geometry is tetrahedral

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Inorganic compounds have coordination number of two, three, four, five, seven and more are very common. The common coordination geometry is an octahedral arrangement around the central atom as show TiF4-

Number of element

Organic compound has hydrocarbon compound and their derivates; H, N, S, O, P, X and may be Sc and Mg.Inorganic compound is the chemistry of everything else, and the elements and their

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everything else, and the elements and their compounds except hydrocarbon and their derivates

PROBLEMS

1. Compare the number of element between organic and inorganic compound

2. Compare the number of bond, geometry of compound between organic and inorganic

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compound between organic and inorganic chemistry

3. Is inorganic chemistry contrast with organic chemistry? Explain your reason!

THE HYSTORY OF INORGANIC CHEMISTRY

Even before alchemy became a subject of study, many reactions were used and the product applied in daily life.Example:• First metal used were probabli gold, copper that can be found in metallic state• In 3000 B.C. silver, tin, antimony, lead were known

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• In 3000 B.C. silver, tin, antimony, lead were known•In 1500 B.C. Iron appear in classical great arroundmediteraninan also four colored glass, ceramic glass•At the first centuries AP, chemistry were active in chine and egypt. The triad to “tranmute” nbase metal into gold •1500 AD chemistry reappeared in europe

•1600 AD, chemistry appeared in art•Roger Bacon (114-194) recognized as one of the first great experimental scientist•By the 17th centuries: found the common strong acid (Chloride acid, Sulfuric acid, and nitric acid)•1869, Becquerel discovered radioactivity •1913 Bohr atomic theory•1926 quantum mechanics of Schrodinger on Heisenberg•1940s, a great expansion of inorganic chemistry

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•1940s, a great expansion of inorganic chemistry•19050s describe the spectra of metal ion, CFT and LFT in coordination compound•1955 discovered organometallic compound

GENESIS OF THE ELEMENT (THE BIG BANG THEORY) AND FORMATION OF THE EARTH

According to the big bang theory, the universe began about 1.8 X lo1' years ago with an extreme concentration of energy in a very small space. In fact, extrapolation back to the time of origin requires zero volume and infinite temperature. Whether this is true

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volume and infinite temperature. Whether this is true or not is still a source of argument, What is almost universally agreed on is that the universe is expanding rapidly, from an initial event during which neutrons were formed and decayed quickly (half-life = 11.3 min) into protons, electrons, and antineutrinos:

n p + e

Lecture IIATOMIC STRUCTURE

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ATOMIC STRUCTURE

Development of atomic model �John Dalton theory � Model atom of

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� Model atom of Thomson�Model atom Rutherford�Model atom of NielsBohr�Model of mechanic wave

JOHN DALTON

The ultimate particle of homogeneous bodies are pereftly alite in weight figure etc. In other word every particle of water is like every other paticle of water, every particle of hydrogen is like every other paticle of hydrogen, etc.

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Atom of differen element has different weight, volume, and propertes

MODEL ATOM OF THOMSON

An atom is a uniform sphere of positive electricity with a radius about 10-8 cm with the electrons embedded in this sphere in such way to give the must stable electrostatic arrangement.

MODEL ATOM OF RUTHERFORD

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MODEL ATOM OF RUTHERFORD

From the experiment of a small fraction of alpha particles where deflected a large angels on passing trough god foil, while most of them passed directly through. It means atom has much empty space and heavy, tiny nucleus carrying a positive charge. Electrons goes around a nucleus in far distance

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As the electrons drop from level nh to nl (h for higher level, 1 for lower level), energy is released in the form of electromagnetic radiation.Conversely, if radiation of the correct

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energy is absorbed by an atom, electrons areraised from level nl to level nh. The inverse-square dependence of energy on nl resultsin energy levels that are far apart in energy at small nl and become much

Parallel discoveries in atomic spectra showed that each element emits light of specific energies when excited by an electric discharge or heat. In 1885, Balmer showed that the energies of visible light emitted by the hydrogen atom are given by the equation

E = RH (1/ni1

Where nh = integer, with nh > 2

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Where nh = integer, with nh > 2RH = Rydberg constant for hydrogen = 1.097 X lo7 m-' = 2.179 X 10-18J and the energy is related to the wavelength, frequency, and wave number of the light, as given by the equation

where h = Planck's constant = 6.626 X J sv = frequency of the light, in s-Ic = speed of light = 2.998 X 10' m s-'h = wavelength of the light, frequently in nmv = wave number of the light, usually in cm-I

E =hv = hc/λ = hcv

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v = wave number of the light, usually in cm-I

The Balmer equation was later made more general, asspectral lines in the ultraviolet and infrared regions ofthe spectrum were discovered, by replacing 22 by nf,with thecondition that nl < nh . These quantities, ni, are calledquantumquantumquantumquantum numbersnumbersnumbersnumbers....

MODEL ATOM OF NIELSBOHR

This theory assumed that negative electrons in atoms move in stable circular orbits around the positive nucleus with no absorption or emission of energy. However, electrons may absorb light of certain specific energies and be excited to orbits of higher energy; they may also emit light of specific energies

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energy; they may also emit light of specific energies and fall to orbits of lower energy.

Model atom of Dalton, Thomson, Rutherford, nielsbohr and mechanic wave

Lecture 3MODEL ATOM OF MECHANIC

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MODEL ATOM OF MECHANIC WAVE

When applied to hydrogen, Bohr's theory worked well; whenatoms with more electrons were considered, the theoryfailed. Complications such as elliptical rather than circularorbits were introduced in an attempt to fit the data to Bohr'stheory.' The developing experimental science of atomicspectroscopy provided extensive data for testing of the

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spectroscopy provided extensive data for testing of theBohr theory and its modifications and forced the theorists towork hard to explain the spectroscopists' observations. Inspite of their efforts, the Bohr theory eventually provedunsatisfactory; the energy levels shown in Figure are validonly for the hydrogen atom. An important characteristic ofthe electron, its wave nature, still needed to be considered.

According to the de Broglie equation,12 proposed in the 1920s, all moving particles have wave properties described by the equation

ƛ = h / mv

where h = wavelength of the particle

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h = wavelength of the particleh = Planck's constantm = mass of the particlev - velocity of thc particle

Particles massive enough to be visible have very shortavelengths, too small to be measured. Electrons, on theother hand, have wave properties because of their verysmall mass.

Electrons moving in circles around the nucleus, as inBohr's theory, can be thought of as forming standingwaves that can be described by the de Broglie equation.However, we no longer believe that it is possible todescribe the motion of an electron in an atom soprecisely. This is a consequence of anotherfundamental principle of modern physics, Heisenberg'sHeisenberg'sHeisenberg'sHeisenberg'suncertaintyuncertaintyuncertaintyuncertainty principle, which states that there is a

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uncertaintyuncertaintyuncertaintyuncertainty principle, which states that there is arelationship between the inherent uncertainties in thelocation and momentum of an electron movingin the x direction:

x px ≥ h / 4Пwherex= uncertainty in the position of the electronp, = uncertainty in the momentum of the electron

The energy of spectra lines can be measured with greatprecision , in turn allowing precise determination of theenergy of electrons in atoms. This precision in energyalso implies precision in momentum (Ap, is small);therefore, according to Heisenberg, there is a largeuncertainty in the location of the electron ( x is large).These concepts mean that we cannot treat electrons assimple particles with their motion described precisely, but

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we must instead consider the wave properties ofelectrons, characterized by a degree of uncertainty intheir location. In other words, instead of being able todescribe precise orbits of electrons, as in the Bohr theory,we can only describe orbitals, regions that describe theprobable location of electrons. The probability of findingthe electron at a particular point in space (also called theelectron density) can be calculated, at least in principle.

The Schrodinger equation describes the wave properties of an electron in terms of its position, mass, total energy, and potential energy. The equation is based on the wavefunction, 9, which 9, which 9, which 9, which describes an electron wave in space; in other describes an electron wave in space; in other describes an electron wave in space; in other describes an electron wave in space; in other words, it describes anwords, it describes anwords, it describes anwords, it describes anatomic orbital. In its simplest

THE SCHRODINGER EQUATION

Ψ=Ψ EH

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notation, the equation is

WhereH = the Hamiltonian operatorE = energy of the electron

= the wave function

Ψ

Ψ=Ψ EH

The Hamiltonian operator (frequently justcalled the Hamiltonian) includes derivativesthat operate on the wave function." Whenthe Hamiltonian is carried out, the result isa constant (the energy) times .... Theoperation can be performed onononon anyanyanyany wavewavewavewave

Ψ

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function describing an atomic orbital.Different orbitals have differentfunctions and different values of E. This isanother way of describing quantization inthat each orbital, characterized by its ownfunction, has a characteristic energy.

Because every psi matches an atomicorbital, there is no limit to the number ofsolutions of the Schrodinger equation for anatom. Each psi describes the waveproperties of a given electron in aparticular orbital. The probability of finding

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particular orbital. The probability of findingan electron at a given point in space isproportional to psi2. A number of conditionsare required for a physically realisticsolution for psy .

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A more detailed look at the Schrodinger equation shows the A more detailed look at the Schrodinger equation shows the A more detailed look at the Schrodinger equation shows the A more detailed look at the Schrodinger equation shows the mathematical originmathematical originmathematical originmathematical originof atomic orbitals. In three dimensions, T may be expressed in T may be expressed in T may be expressed in T may be expressed in terms of Cartesian coordinatesterms of Cartesian coordinatesterms of Cartesian coordinatesterms of Cartesian coordinates( x , y, z ) or in terms of spherical coordinates (r, 0, +). Spherical coordinates,as shown in , are especially useful in that r represents the distance from the

The Schrodinger equation The Schrodinger equation The Schrodinger equation The Schrodinger equation

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distance from thenucleus. The spherical coordinate 0 is the angle from the z axis, varying from 0 to n ,and 4 is the angle from the x axis, varying from 0 to 2n. It is possible to convert betweenCartesian and spherical coordinates using the following expressions:

x = r sin 0 cos +y = r sin 0 sin +z = r cos 0

The angular functionsThe angular functionsThe angular functionsThe angular functionsThe angular functions φ and θ determine how the probability determine how the probability determine how the probability determine how the probability changes from point to changes from point to changes from point to changes from point to point at a given distance from the center of the atom; in other words, they give the shape of the orbitals and their orientation in space. The angular functions φ and θ are determined

by the quantum numbers 1 and ml. The shapes of s, p, and d orbitals are given.

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The radial functionsThe radial functionsThe radial functionsThe radial functionsThe radial factor R(r) is determined by the quantum numbers n and 1, the principal and angular momentum quantum numbers.This function describes the probability of finding the electron at a given distance from the nucleus, summed over all angles, with the 4πr2 factor the result of integrating over all angles. The radial wave functions and radial probability functions are plotted for the n = 1, 2, and 3 orbitals

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S,p,d,orbitals

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LECTURE 4Electron configuration

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Electron configuration

THETHETHETHE AUFBAUAUFBAUAUFBAUAUFBAU PRINCIPLEPRINCIPLEPRINCIPLEPRINCIPLELimitations on the values of the quantum numbers lead tothe familiar AufbauAufbauAufbauAufbau (German,(German,(German,(German, Auflau, building up)principle,principle,principle,principle, wherewherewherewhere thethethethe buildupbuildupbuildupbuildup ofofofof electronselectronselectronselectrons inininin atomsatomsatomsatoms resultsresultsresultsresultsfromfromfromfrom continually increasing the quantum numbers. Anycombination of the quantum numbers presented so farcorrectly describes electron behavior in a hydrogen

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correctly describes electron behavior in a hydrogenatom, where there is only one electron. However,interactions between electrons in polyelectronic atomsrequire that the order of filling of orbitals be specifiedwhen more than one electron is in the same atom. In thisprocess, we start with the lowest n, 1, and ml,ml,ml,ml, valuesvaluesvaluesvalues ((((1111,,,,0000,,,, andandandand 1 0, respectively) and either of the rnrnrnrn,,,, valuesvaluesvaluesvalues (we(we(we(wewillwillwillwill arbitrarilyarbitrarilyarbitrarilyarbitrarily useuseuseuse ---- 3333 first)first)first)first).... ThreeThreeThreeThree rulesrulesrulesrules will then give usthe proper order for the remaining electrons as weincrease the quantum numbers in the order ml,ml,ml,ml, m,,m,,m,,m,, IIII ,,,, andandandandnnnn....

1. Electrons are placed in orbitalsto give the lowest total energy tothe atom. This means that thelowest values of n and I are filled

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lowest values of n and I are filledfirst. Because the orbitals withineach set (p,(p,(p,(p, d,d,d,d, etcetcetcetc....)))) havehavehavehave thethethethe samesamesamesameenergy,energy,energy,energy, thethethethe ordersordersordersorders forforforfor valuesvaluesvaluesvalues ofofofof mlmlmlmlandandandand m,m,m,m, areareareareindeterminate.

Energy level fill in electron configuration

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The Pauli exclusion principle The Pauli exclusion principle The Pauli exclusion principle The Pauli exclusion principle requires that each requires that each requires that each requires that each electron in an atom have a unique set of quantum numbers. At least one quantum

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numbers. At least one quantum number must be different from those of every other electron. This principle does not come from the Schrodinger equation, but from experimental determination of electronic structures.

Hund'sHund'sHund'sHund's rule of maximum multiplicity requires that electrons be rule of maximum multiplicity requires that electrons be rule of maximum multiplicity requires that electrons be rule of maximum multiplicity requires that electrons be placed in placed in placed in placed in orbitalsorbitalsorbitalsorbitals so as to give the maximum total spin possible (or the maximum number of parallel spins). Two electrons in the same orbital have a higher energy than two electrons in different orbitals, caused by electrostatic repulsion (electrons in the same orbital repel each other more than electrons in separate orbitals). Therefore, this rule is a consequence of the lowest possible energy rule (Rule 1). When there are one to

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lowest possible energy rule (Rule 1). When there are one to six electrons in p orbitals, the required arrangements are those given table below. The multiplicity is the number of unpaired . The multiplicity is the number of unpaired . The multiplicity is the number of unpaired . The multiplicity is the number of unpaired electrons plus 1, or electrons plus 1, or electrons plus 1, or electrons plus 1, or n + I. n + I. n + I. n + I. This is the number of possible energy levels that depend on the orientation of the net magnetic moment in a magnetic field. Any other arrangement of electrons results in fewer unpaired electrons.

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This rule is a consequence of the energy required for pairing electrons in thesame orbital. When two electrons occupy the same part of the space around an atom,they repel each other because of their mutual negative charges with a Coulombic energyCoulombic energyCoulombic energyCoulombic energyof repulsion, II,, per pair of electrons. As a result, this repulsive of repulsion, II,, per pair of electrons. As a result, this repulsive of repulsion, II,, per pair of electrons. As a result, this repulsive of repulsion, II,, per pair of electrons. As a result, this repulsive force favorsforce favorsforce favorsforce favorselectrons in different orbitals (different regions of space) over

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electrons in different orbitals (different regions of space) over electrons in the sameorbitals.In addition, there is an exchange energy, II,, which arises from exchange energy, II,, which arises from exchange energy, II,, which arises from exchange energy, II,, which arises from purely quantumpurely quantumpurely quantumpurely quantummechanical considerations. This energy depends on the number of possible exchangesbetween two electrons with the same energy and the same spin.For example, the electron configuration of a carbon atom is 1s1s1s1s2222

2s2s2s2s2222 2p2p2p2p2 2 2 2 three arrangement of the 2p electrons can be three arrangement of the 2p electrons can be three arrangement of the 2p electrons can be three arrangement of the 2p electrons can be

In the first two cases there is only one possible way to arrange the electrons to give the same diagram, because there is only a single electron in each having + or spin. However, in the third case there are two possible ways in which the electrons can be arranged:

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The exchange energy is IIe, per possible exchange of parallel electrons and is negative. The higher the number of possible exchanges, the lower the energy. Consequently, the third configuration is lower in energy than the second by IIe

The results may be summarized in an energy diagram:

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These two pairing terms add to produce the total pairing energy, II:

II= IIc + IIeThe Coulombic energy, II,, is positive and is nearly constant for each pair of

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1. a. Find the possible values for the 1. a. Find the possible values for the 1. a. Find the possible values for the 1. a. Find the possible values for the I and I and I and I and rnlrnlrnlrnlquantum quantum quantum quantum numbers for a 5d electron, a 4fnumbers for a 5d electron, a 4fnumbers for a 5d electron, a 4fnumbers for a 5d electron, a 4f

electron, and a 7g electron.b. Find the possible values for all four quantum b. Find the possible values for all four quantum b. Find the possible values for all four quantum b. Find the possible values for all four quantum

numbers for a 3d electron.numbers for a 3d electron.numbers for a 3d electron.numbers for a 3d electron.2. Give explanations of the following phenomena:2. Give explanations of the following phenomena:2. Give explanations of the following phenomena:2. Give explanations of the following phenomena:

a. The a. The a. The a. The electron configuration of Cr is [ configuration of Cr is [ configuration of Cr is [ configuration of Cr is [ ArArArAr ] 4 s ] 4 s ] 4 s ] 4 s 1111 3 3 3 3

Problems

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a. The a. The a. The a. The electron configuration of Cr is [ configuration of Cr is [ configuration of Cr is [ configuration of Cr is [ ArArArAr ] 4 s ] 4 s ] 4 s ] 4 s 1111 3 3 3 3 dddd5555 rather than [ rather than [ rather than [ rather than [ ArArArAr ] 4 s ] 4 s ] 4 s ] 4 s 22223 d3 d3 d3 d4444....

b. The electron configuration of Ti is [ Ar ] 4 s 4 s 4 s 4 s 2222

3 d3 d3 d3 d2 2 2 2 but that of Cr2 +is [ [ [ [ ArArArAr ] 3d] 3d] 3d] 3d4444....3. Give electron configurations for the following:3. Give electron configurations for the following:3. Give electron configurations for the following:3. Give electron configurations for the following:

a. Sc b. Br c. Fea. Sc b. Br c. Fea. Sc b. Br c. Fea. Sc b. Br c. Fe3+3+3+3+ d. Bi e. Sbd. Bi e. Sbd. Bi e. Sbd. Bi e. Sb4. Which 2+ ion has five 3d electrons? Which one 4. Which 2+ ion has five 3d electrons? Which one 4. Which 2+ ion has five 3d electrons? Which one 4. Which 2+ ion has five 3d electrons? Which one has two 3d electrons?has two 3d electrons?has two 3d electrons?has two 3d electrons?

Lecture 5The periodic

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The periodic table

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One of the simplest methods that fits most atoms uses the periodic table blocked out as in Figure before. The electron configurations of hydrogen and helium are clearly 1s'and 1s2. After that, the elements in the first two columns on the left (Groups 1 and 2 or IA and IIA) are filling s orbitals, with 1 = 0; those in the six columns on the right (Groups 13 to

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18 or IIIA to VIIIA) are filling p orbitals, with 1 = 1; and the ten in themiddle (the transition elements, Groups 3 to 12 or IIIB to IIB) are filling d orbitals, with 1 = 2. The lanthanide and actinide series (numbers 58 to 71 and 90 to 103) are filling f orbitals, with 1 = 3. Either of these two methods is too simple, as shown in the following paragraphs, but they do fit most atoms and

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SHIELDINGSHIELDINGSHIELDINGSHIELDINGIn atoms with more than one electron, energies of specific levels are difficult to predict quantitatively, but one of the more common approaches is to use the idea of shielding. Each electron acts as a shield for electrons farther out from the nucleus, reducing the attraction between the nucleus and the distant electrons. Although the quantum number n is most important in determining the energy, 1 must also be included in the calculation of the energy in atoms with more than one

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in the calculation of the energy in atoms with more than one electron. As the atomic number increases, the electrons are drawn toward the nucleus and the orbital energies become more negative. Although the energies decrease with increasing Z, the changes are irregular because of shielding of outer electrons by inner electrons. The resulting order of orbital filling for the electrons

As a result of shielding and other more subtle interactions between the electrons, the simple order of orbitals (in order of energy increasing with increasing n) holds only at very low atomic number Z and for the innermost electrons of any atom. For the outer orbitals, the increasing energy difference between levels with the same n but different 1 values forces the overlap of energy levels with n = 3

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values forces the overlap of energy levels with n = 3 and n = 4, and 4s fills before 3d. In a similar fashion, 5s fills before 4d, 6s before Sd, 4f before Sd, and 5f before 6d later formulated a set of simple rules that serve as a rough guide to this effect.He defined the effective nuclear charge Z* as a measure of the nuclear attraction for an electron. Z* can be calculated from Z* = Z - S, where Z is the nuclear charge and S I the shielding constant. The

1. The electronic structure of the atom is written in groupings as follows: (1s)(2s,2 p) (3s,3 p) ( 3 4 ( 4s,4 p) ( 4 4 ( 4 f ) ( 5s,5 ~ 1e,tc .2. Electrons in higher groups (to the right in the list above) do not shield those in .lower groups.3. For ns or np valence electrons:a. Electrons in the same ns, np group contribute 0.35, except the Is, where 0.30

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the Is, where 0.30works better.b. Electrons in the n - 1 group contribute 0.85.c. Electrons in the n - 2 or lower groups contribute 1.00.4. For nd and nf valence electrons:a. Electrons in the same nd or nf group contribute 0.35.b. Electrons in groups to the left contribute 1.00.The shielding constant S obtained from the sum of the contributions above is subtracted from the nuclear charge Z to obtain the effective nuclear charge Z* affecting the selected electron.

Justification for Slater's rules (aside from the fact that they work) comes from theelectron probability curves for the orbitals. The s and p orbitals have higher probabilitiesnear the nucleus than do d orbitals of the same n, as shown earlier in Figure 2-7.Therefore, the shielding of 3d electrons by (3s, 3p) electrons is calculated as 100% effective(a contribution of 1.00). At the same time, shielding of 3s or 3p electrons by(2s, 2p) electrons is only 85% effective (a contribution of0.85), because the

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(2s, 2p) electrons is only 85% effective (a contribution of0.85), because the 3s and 3porbitals have regions of significant probability close to the nucleus. Therefore, electronsin these orbitals are not completely shielded by (2s, 2p) electrons.A complication arises at Cr ( 2 = 24) and Cu ( 2 = 29) in the first transition seriesand in an increasing number of atoms under them in the second and third transitionseries. This effect places an extra electron in the 3d level and removes one electron fromthe 4s level. Cr, for example, has a configuration of [ A r ] 4 s1 3 d 5(rather

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As the nuclear charge increases, the electrons are more strongly attracted and the energy levels decrease in energy, becoming more stable, with the d orbitals changing more rapidly

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than the s orbitalsbecause the d orbitals are not shielded as well from the nucleus. Electrons fill the lowest available orbitals in order up to their capacity.

LECTURE 6LECTURE 6LECTURE 6LECTURE 6PERIODIC PROPERTIES OF ATOMSPERIODIC PROPERTIES OF ATOMSPERIODIC PROPERTIES OF ATOMSPERIODIC PROPERTIES OF ATOMS

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PERIODIC PROPERTIES OF ATOMSPERIODIC PROPERTIES OF ATOMSPERIODIC PROPERTIES OF ATOMSPERIODIC PROPERTIES OF ATOMS

IONIZATION ENERGYIONIZATION ENERGYIONIZATION ENERGYIONIZATION ENERGYThe ionization energy, also known as the ionization potential, is the energy required to remove an electron from a gaseous atom or ion:

An+ (g) A (n+1)+ (g) + e ionization energy = AUAUAUAU

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energy = AUAUAUAUwhere n = 1 (first ionization energy), 1, 2, (second, third, . . . ) As would be expected from the effects of shielding, the ionization energy varies with different nuclei and different numbers of electrons. Trends for the first ionization energies of the early elements in the periodic table are shown in Figure below. The general trend across a period is an increase in ionization energy as the nuclear charge increases

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ELECTRON AFFINITYELECTRON AFFINITYELECTRON AFFINITYELECTRON AFFINITYElectron affinity can be defined as the energy required to remove an electron from anegative ion:A- (g) A (g) + e ionization energy = AU (or AU (or AU (or AU (or EA)EA)EA)EA)(Historically, the definition is -AU for the reverse reaction, adding an electron to theneutral atom. The definition we use avoids the sign change.) Because of the similarity

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the similarityof this reaction to the ionization for an atom, electron affinity is sometimes described asthe zero ionization energy. This reaction is endothermic (positive AU), except for thenoble gases and the alkaline earth elements. The pattern of electron affinities withchanging Z shown in Figure above is similar to that of the ionization energies, but for onelarger Z value (one more electron for each species) and with much smaller absolutenumbers. For either of the reactions, removal of the first electron past a

COVALENT AND IONIC RADIICOVALENT AND IONIC RADIICOVALENT AND IONIC RADIICOVALENT AND IONIC RADIIThc sizes of atoms and ions are also related to the ionization energies and electron affinities.As the nuclear charge increases, the electrons are pulled in toward the center of the atom, and the size of any particular orbital decreases. On the other hand, as the nuclear charge increases, more electrons are added to the atom and their mutual repulsion keeps the outer orbitals large. The interaction of these two effects (increasing nuclear charge and increasing number of electrons) results in a gradual decrease in atomic size across each period. Following table gives nonpolar covalent radii, calculated for ideal molecules with no polarity.

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There are similar problems in determining the size of ions. Because the stable ions of the different elements have different charges and different numbers of electrons, as well as different crystal structures for their compounds, it is difficult to find a suitable set of numbers for comparison. Earlier data were based on Pauling's approach, in which the ratio of the radii of isoelectronic ions was assumed to be equal to the ratio of their effective nuclear charges More recent calculations are based on a number of considerations, including electron density maps from X-ray data that show larger cations and smaller anions than those previously found.

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The values in folowing Table show that anions are generally larger lhan cations with similar numbers of electrons (F -and Na+ differ only in nuclear charge, but the radius of fluoride is 37% larger). The radius decreases as nuclear charge increases for ions with the same electronic structure, such as 02-,F -, N a +and Mg+2, with a much larger change with nuclear charge for the cations. Within a family, the ionic

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charge for the cations. Within a family, the ionic radius increases as Z increases because of the larger number of electrons in the ions and, for the same element,the radius decreases with increasing charge on the cation. Examples of thesetrends are shown in those following Tables

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problems

1. Using Slater's rules, determine Z* for1. Using Slater's rules, determine Z* for1. Using Slater's rules, determine Z* for1. Using Slater's rules, determine Z* fora. A 3p electron in P, S, C1, and Ar. Is the calculated value a. A 3p electron in P, S, C1, and Ar. Is the calculated value a. A 3p electron in P, S, C1, and Ar. Is the calculated value a. A 3p electron in P, S, C1, and Ar. Is the calculated value

of Z* consistent with theof Z* consistent with theof Z* consistent with theof Z* consistent with therelative sizes of these atoms?

b. A 2p electron in 0-2 F-1, Na +1 ,and Mg+2. Is the calculated value of Z* consistent

with the relative sizes of these ions?c. A 4s and a 3d electron of Cu. Which type of electron is

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c. A 4s and a 3d electron of Cu. Which type of electron is more likely to be lost when

copper forms a positive ion?d. A 4 f electron in Ce, Pr, and Nd. There is a decrease in

size, commonly known as thelanthanide contraction, with increasing atomic number

in the lanthanides. Areyour values of Z* consistent with this trend?

2.Select the better choice in each of the following, and explain your selection briefly.

Lecture 7

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The Ionic Compound

• Simple ionic compound form only between very active metallic elements and very active nonmetals. Very active metallic elements form the cation ( positive ion ), that loss electrons, and the active nonmetals gain electron to form anion ( negative ion ). The attraction between posititve and

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anion ( negative ion ). The attraction between posititve and negative ion result in an ionic bond.

• Two importan requisites are that

- the ionization energy to form cation and

- electron affinity to form anion

The requirment for ionic bonding;

• 1) The atoms of one element must be able to lose one or two ( rarely three) electrons without undue energy input.

• 2) The atom of other element must be able to accept one or two electrons ( almost never three) without undue energy input.

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energy input.

The ristrict ionic bonding to compounds between active metals group IA, IIA, IIIA and some lower state of transition metals (formins cations) – and the most active non metals (group VII A and VI A and VA) (forming anions)

Formation of ionic bonds is illustrated for sodium chloride by :

Na (s) Na (g) Na + (g)

½ Cl 2 (g) Cl (g) Cl –(g)

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½ Cl 2 (g) Cl (g) Cl (g)

Na + (g) + Cl –(g) NaCl (s)

Where Na + and Cl – both have filled octets

The energy of interactions of two charged particles is given by

E = q + q –

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E = q q

4 π r €0

q + and q – arethe charges, r is distance of separation € is the permittif of the medium

The formation of sodium chloride not simple since the real sodium chloride in Na+ and CL- in the formation of lattice in the face center cubic structure. The coordination number(C.N) of

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structure. The coordination number(C.N) of both ions in the sodium chloride lattice is 6. there are six chloride ions about each sodium ion and six sodium ions about each chloride ion.

Interaction between Na+ and CL- in the lattice state is bigger than those in gas state, because the lattice has a lattice energy.

LATTICE ENERGY

The energy of the crystal lattice of an ionic compound is the energy released when ions (gas state) come together from infinite a crystal

M + (g) + X – (g) MX (s)

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(g) (g) (s)

It may be treated adequately by a simple electro static model. Although we shall include non electrostatic energies , such as the repulsion of closed cell and more sophisticated treatment include such factors as dispersion forces and zero energy, simple electrostatic accounts for about 90 % of the bonding energy.

the theoretical treatment of ionic lattice energy was initiated by Born and Landle. They derivates from Coulombs law.

E = Z + Z –

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E = Z Z

4 π r €0

Now the crystals lattice there will be more interactions than the simple ne in an ion par. In the sodium chloride lattice, there attraction to the six nearest neighbor of opposite charge and repulsion by the twelve neighbors of live charge, etc . so in the pair of ion equation in lattice become

E = A Z + Z – e2

4 π r €0

A = Madelung constantA = 6- 12/√2 + 8√3 - …….

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Born suggested repulsion energyEr = B/rn

NA = Avogadro numberThe total lattice energy

U = A NA Z + Z – e2 (1- 1/n)4 π r €0

BORN HABER CYLE

His law state that enthalpy of a reaction is the same whether the reaction take place. In one or several steps. It is necessary consequence of the first law of thermodynamics concerning the conservations energy. This is a Born Haber cycle, that apply the Hass law to the enthalpy of formation of an ionic solid.

Hf = S + D + I + A + US = Enthalpy of atomization or enthalpy of sublimation of the

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S = Enthalpy of atomization or enthalpy of sublimation of the metal

D = enthalpy of dissociation of the atomic molecule or Enthalpy of atomization.

I = energy of ionization A = electron affinityU = lattice energy

Lecture 8Ionic radii and properties of ionic

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Ionic radii and properties of ionic compound

The distance between two ionic compound (d) the same with the sum of ionic radii of cation (rA+ ) and that of anion ( rB- )

d= rA+ + rB-

A+ B-

rA+ rB-

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If it possible to measure the distance (d) with high degree of accuracy , experimentally d= rA+ + rB- And could not tell where each .

rA+ rB-

We know only that Cation ends and the anion begins.X-ray crystallographer determine the structure of compound such as NaCl,

usually only the spacing of ions is determined ( d is determined).d= rA+ + rB-

Example: d KF = 2,66 A0 ∆r= 0,35d NaF = 2,31 A0 d Kf- d NaF = 0,3 rK+ + rNa+ =0,35

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d KCl = 3,14 A0

d NaCl = 2,81 A0 d KCl- d NaCl = rK+ + rNa+ = 0,33

Conclusion: r ionic compound is constant.

Pauling Approach about Ionic Radii

The ratio of the radii of isoelectronic ions was assumed reverse to the ratio of their effective nuclear charge.Na+ = 1s2 , 2s2 ,2p6

F- = 1s2 , 2s2 ,2p6

Z* = Z – S

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Z* = Z – SZ* Na+ = 11-4,5= 6,85Z* F- = 9- 4,5 = 4,85r Na+ / r F- = Z* F- / Z* Na+

r Na+ / r F- = 4,85/6,85 , r Na+ / r F- = 0,71r Na+ = 0,71 r F- …………(1)

r Na+ + r F- = 2,31…………..(2)

r Na+ = 0,96

r F- = 1,35

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Problem

Calculate the r K+ and rCl - if the d KCl= 3,14!

Properties of Ionic Compound

1) Streochemistry

2) Melting and boiling point

3) Hardness

4) Solubility

5) Conductivity

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5) Conductivity

6) Brittleness

1) Streochemistry

Ionic bond are quite strong and are omnidirectional. Ionic forcesextend throughout the space and equally strong in all direction.

2)Melting and boiling point

Melting point and boiling point of ionic compound is high, doe to

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Melting point and boiling point of ionic compound is high, doe tostrong electronic union and cation interactions, extending throughcrystal lattice. The closer the ion the stronger will be the attractionforces and the higher will be the melting point.

3) Hardness

The multivalent electrostatic attraction forces of ionic crystal makethe crystal hard. Hardness increase with the decriasing interionicdistance, and increase the ionic charge.

4)Solubility

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4)Solubility

Ionic compound dissolve in polar solvent having high dielectricconstant, doe to :

� Dicteace in the interaction force among ion the a electric medium.

� Ion solvent dipole interaction.

5) Conductivity

� In the solid state, ionic compound have extremely low conductanceas the ion are the held tightly.

� In molten and solution stall, ionic compounds conduct electricity,because the formation of free mobile ion..

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6) Brittleness

If safficient energy is supplied to the crystal layer of a unit cell, theattraction force become repulsive do to the anion-anion and cation-cation repulsion and the crystal crumbles. The ionic crystal hard butbrittle, can be powdered.

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94

Covalent Bonding

Week 1


95“Read in the name your God who created “ )I( Inorganic Chemistry 3 )I( Miftahul Khair, M.Sc ([email protected])

Intro

• Miftahul Khair

• Room : CF 099 Chem Dept UNP

• [email protected]

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• Check regularly your portal for taking the ppt slides, course materials, pre exam problems, etc (if any)

96

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Models of Covalent Bonding

The properties of atoms � atomic orbitals, so the properties of covalent compounds � molecular orbitals.

An electron in a molecular orbital is the property of the whole molecule, not of an individual atom.

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Molecular orbital energy levels : - aspects of chemical bonding that are difficult to comprehend in terms of the simple Lewis electron-dot representations.

Michael Faraday: oxygen gas is attracted into a magnetic field � paramagnetic � unpaired electrons unpaired electrons unpaired electrons unpaired electrons Later, bond strength studies showed that the O2 molecule has a double bonddouble bonddouble bonddouble bond.

the ground-state molecular

electron-dot diagrams

OOOO2222 molecule molecule molecule molecule

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orbital diagram

Introduction to Molecular Orbitals

When two atoms approach each other, their atomic orbitals overlap. The electrons no longer belong to one atom but to the molecule as a whole.

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1. For orbitals to overlap, the signs on the overlapping lobes must be the same.1. For orbitals to overlap, the signs on the overlapping lobes must be the same.1. For orbitals to overlap, the signs on the overlapping lobes must be the same.1. For orbitals to overlap, the signs on the overlapping lobes must be the same.2. Whenever two atomic orbitals mix, two molecular orbitals are formed, : 2. Whenever two atomic orbitals mix, two molecular orbitals are formed, : 2. Whenever two atomic orbitals mix, two molecular orbitals are formed, : 2. Whenever two atomic orbitals mix, two molecular orbitals are formed, : bonding and antibonding. 3. For significant mixing to occur, the atomic orbitals must be of similar energy.3. For significant mixing to occur, the atomic orbitals must be of similar energy.3. For significant mixing to occur, the atomic orbitals must be of similar energy.3. For significant mixing to occur, the atomic orbitals must be of similar energy.

Statements about molecular orbitals:Statements about molecular orbitals:Statements about molecular orbitals:Statements about molecular orbitals:

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3. For significant mixing to occur, the atomic orbitals must be of similar energy.3. For significant mixing to occur, the atomic orbitals must be of similar energy.3. For significant mixing to occur, the atomic orbitals must be of similar energy.3. For significant mixing to occur, the atomic orbitals must be of similar energy.4. Each molecular orbital can hold a maximum of two electrons, one with 4. Each molecular orbital can hold a maximum of two electrons, one with 4. Each molecular orbital can hold a maximum of two electrons, one with 4. Each molecular orbital can hold a maximum of two electrons, one with spin +1/2 , the other -1/25. The electron configuration of a molecule can be constructed by using the 5. The electron configuration of a molecule can be constructed by using the 5. The electron configuration of a molecule can be constructed by using the 5. The electron configuration of a molecule can be constructed by using the Aufbau principle by filling the lowest energy molecular orbitals in sequence.6. When electrons are placed in different molecular orbitals of equal energy, 6. When electrons are placed in different molecular orbitals of equal energy, 6. When electrons are placed in different molecular orbitals of equal energy, 6. When electrons are placed in different molecular orbitals of equal energy, the parallel arrangement (Hund’s rule) will have the lowest energy.7. The bond order in a diatomic molecule is defined as the number of bonding7. The bond order in a diatomic molecule is defined as the number of bonding7. The bond order in a diatomic molecule is defined as the number of bonding7. The bond order in a diatomic molecule is defined as the number of bondingelectron pairs minus the number of antibonding pairs.

Molecular Orbitals for Period 1 Diatomic Molecules

Start with the simplest diatomic species H2+

H + H+ � H2+ molecular ion.

electron configuration: (σ1s)1

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Simultaneous attraction of the electron to

two H nuclei � Energy of the electron σ1s<

1s

•net reduction in total electron energy is the

driving force in covalent bond formation.


Total

e-

Bond

order

bond

length (pm)

bond strength

(kJ/mol)

1 1/2 106 255

H + H � H2

Hydrogen molecule, HHydrogen molecule, HHydrogen molecule, HHydrogen molecule, H2222

The greater the bond order, the greater the strength of the bond and the shorter the bond length

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Total

e-

Bond

order

bond

length (pm)

bond strength

(kJ/mol)

2 1 74 436

Match!

He + He+ � He2+ molecular ion.

He2+ molecular ion.

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electron configuration: (σ1s)2 (σ*1s)

1

Total

e-

Bond

order

bond

length (pm)

bond strength

(kJ/mol)

31-1/2=

1/2108 251 Weaker bond is confirmed by the bond

length and bond energy, the same as those of the dihydrogen ion.

He + He � He2 molecule.

He2 molecule

2 e- decrease in energy while, 2e-increase in energy by the same quantity.

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electron configuration: (σ1s)2 (σ*1s)

2

Total

e-

Bond

order

bond

length (pm)

bond strength

(kJ/mol)

4 1-1= 0

increase in energy by the same quantity. � no net decrease in energy

net bond order = 0

orororor

So, no covalent bonding !

Molecular Orbitals for Period 2 Diatomic MoleculesAt a new period, the inner (core) e- become irrelevant to the bonding process.

Li (2s1) and Be (2s2) �only construct a molecular orbital energy diagram corresponding to the 2s atomic orbitals.

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The outermost occupied orbitals /the frontier orbitals � crucial orbitals for bonding.

HOMO/highest occupied molecular orbitals � crucial for bonding LUMO/ lowest unoccupied molecular orbitals � crucial for reaction.

Li + Li � Li2 molecule.

Li2 (g) molecule

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Occupancy of frontier(valence) molecular orbital : (σ2s)

2

Total

e-

Bond

order

bond

length (pm)

bond strength

(kJ/mol)

2 1 OK OK

Try Be2 molecule ! Possible ?

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Heavier Period 2 elements

Formation of molecular orbitals from 2p atomic orbitals. Mix in two ways: 1.1.1.1.end to end end to end end to end end to end � a pair of bonding and antibonding orbitals is formed and resembles those of the σ1s orbitals. � σ2p and σ*2p

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109

resembles those of the σ1s orbitals. � σ2p and σ*2p

σ bond = bond formed by atomic orbital overlap along the axis joining the two nuclear centers.

As orbitals can only overlap if the signs of the lobes are the same—in this case, we show positive to positive.

2. S2. S2. S2. Side to sideide to sideide to sideide to side. � The bonding and antibonding molecular orbitals Π, for both px and py

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110

every atom has 3 2p atomic orbitals, �3 bonding and 3 antibonding

molecular orbitals.

Bonding models � explain experimental observations.

the shorter the bond length and the higher the bond energy, the stronger the bond

Bond energy (kJ/mol) Bond order

200–300 single bonds,

500–600 double bonds

900–1000 triple bonds

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Diatomic molecules, N2, O2, and F2, the bonding and antibonding orbitals formed from both 1s and 2s

atomic orbitals are filled

•No net bonding contribution from these orbitals.

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these orbitals.

•Only consider the filling of the molecular orbitals derived from the

2p atomic orbitals.

the Period 2 elements beyond N2Hund’s rule must works,

• two unpaired e-• bond order

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• bond order

NNNN2222

At the beginning of the period, the

levels differ in energy by only about 0.2

MJ/mol. In these circumstances, the

wave functions for the 2s and 2p

orbitals become mixed.

�increase in energy of the σ2p

molecular orbital to the point where it

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115

electron configuration: (π2p)4 (σ2p)2

molecular orbital to the point where it

has greater energy than the π2p

orbital.

�This ordering of orbitals applies to

dinitrogen and the preceding elements

in Period 2, the σ- π crossover

occurring between N2 and O2

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Molecular Orbitals for Heteronuclear Diatomic Molecules

the same period : Z >, Zeff > � orbital energies <

the molecular orbitals derived primarily from the 2s atomicorbitals of one element overlap significantly in energy with those derived from the 2p atomic

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derived from the 2p atomic orbitals of the other element.

two molecular orbitals whose energiesare between those of the contributing atomic orbitals. These orbitals, σNB, are defi ned as

nonbonding molecular orbitals; that is, they do not contribute significantly to the bonding

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Molecular orbital diagrams can also be constructed to develop bonding schemes for molecules containing more than two atoms.

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But, its better to predict of the shapes of molecules than orbital energy levels.

A Brief Review of Lewis Structures

Lewis; atoms combine by sharing electron pairs

Molecular orbital � bond length and bond energy in covalent molecules.

Lewis � molecular shape of complex molecules.

Lewis (electron-dot) :

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Lewis (electron-dot) :

the driving force of bond formation as being the attainment by each

atom in the molecule of an octet of electrons in its outer (valence)

energy level (except hydrogen) , completion of the octet is

accomplished by a sharing of electron pairs between bonded atoms.

Constructing ElectronConstructing ElectronConstructing ElectronConstructing Electron----Dot DiagramsDot DiagramsDot DiagramsDot Diagrams

1.1.1.1. Identify the central atom, (lower Identify the central atom, (lower Identify the central atom, (lower Identify the central atom, (lower electronegativityelectronegativityelectronegativityelectronegativity)))).Write the symbol of the central atom and place the

symbols of the other atoms around the central atom. 2. Count the total number of valence electrons. 2. Count the total number of valence electrons. 2. Count the total number of valence electrons. 2. Count the total number of valence electrons.

3. Place an electron pair (single covalent bond) between 3. Place an electron pair (single covalent bond) between 3. Place an electron pair (single covalent bond) between 3. Place an electron pair (single covalent bond) between the central atomthe central atomthe central atomthe central atomand each of the surrounding atoms. Add lone pairs to the surrounding atoms. Then any excess electrons are

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the surrounding atoms. Then any excess electrons are added to the central atom.4. If the number of electrons on the central atom is less 4. If the number of electrons on the central atom is less 4. If the number of electrons on the central atom is less 4. If the number of electrons on the central atom is less than eight and there than eight and there than eight and there than eight and there are “leftover” electrons, add lone pairs to the central atom. If the number of electrons on the central atom is less than eight and there are no more electrons, construct double and triple construct double and triple construct double and triple construct double and triple bonds using lone pairs from surrounding atoms.

Exceeding the OctetExceeding the OctetExceeding the OctetExceeding the Octet

< 8 e- or > 8 e- in central atom

the use of d orbitals was invoked to make possible a theoretical maximum of 18 bonding e-

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Partial Bond Order

In some cases, the only structure that can be drawn does not correlate with our measured bond information

Ex: NO3-

N-O = 122 pm. N-Otheoritical = 144 pm.

resonance.

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resonance.

fractional bond order

average bond order =1 1/3.

Formal Charge

In some cases, we can draw more than one feasible electron-dot diagram, one such example being dinitrogen oxide. Try !

Ex: N2O

asymmetric linearwith a central N atom.

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which possibilities are unrealistic� formal charge

formal charge : •divide the bonding e- equally among the constituent atoms and•Compare number of assigned electrons for each atom with its original number of valence e-. •Any difference � charge sign

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the lowest energy structurewill be the one with the smallest formal charges on

the atoms

(c) is eliminated

optimum representation= resonance mixture of a and b.

Valence-Shell Electron-Pair Repulsion Rules

•Electron-dot diagrams can be used to derive the probable

also called the electron domain (ED) model.

repulsions between electron pairs in theoutermost occupied energy levels on a central atom cause those electron pairs to be located as far from each other as

is geometrically possible

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•Electron-dot diagrams can be used to derive the probable molecular shape. •use a very simplistic set of rules that tells us nothing about the bonding : ignore the differences between the energies ofthe s, p, and d orbitals and simply regard them as degenerate

•effective at predicting molecular shapes

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Vocabulary:

“domain”

= any electron pair, or any double or triple bond is considered one domain.

“lone pair” = “non“lone pair” = “non--bonding pair” = “unshared pair”bonding pair” = “unshared pair”


“lone pair” = “non“lone pair” = “non--bonding pair” = “unshared pair”bonding pair” = “unshared pair”

= any electron pair that is not involved in bonding= any electron pair that is not involved in bonding

“bonding pair” = “shared pair”“bonding pair” = “shared pair”

= any electron = any electron pairpair that is involved in bondingthat is involved in bonding

Steps for using VSEPR:

1. Draw a Lewis Dot Structure.

2. Predict the geometry around the central atom.

3. Predict the molecular shape.

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… also, we can try and predict the angles between atoms.… also, we can try and predict the angles between atoms.

All e- pairs push each other as far apart as possible.

• Shared (bonding) pairs are “stretched” between two atoms that want them.

• “Longer & Thinner”

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• Unshared (non-bonding) pairs are not “stretched.”

• “Shorter & Thicker”

2 domains on central atom

LINEAR

• 2 domains

• both are bonding pairs

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• both are bonding pairs

• They push each other to opposite sides of center (180° apart).

BeCl2


TRIGONAL PLANAR

• 3 domains

• all are bonding pairs

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• all are bonding pairs

• They push each other apart equally at 120°degrees.

GaF3


BENT

• 3 domains:• 2 are bonding pairs

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• 2 are bonding pairs

• 1 is a lone pair

• The 2 bonding pairs are pushed apart by 3rd pair (not seen)

SnF2

NOTE:

• The geometry around the central atom is trigonal planar.

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planar.

• The molecular shape is bent.

SnF2


TETRAHEDRAL

• 4 domains

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• Each repels the other equally - 109.5° - not the expected 90°.

• Think in 3D.

CH4

4 e- pairs on central atom

TRIGONAL PYRAMIDAL

• 4 domains

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• 4 domains• 3 bonding pairs

• 1 lone pair

• The thicker, lone pair forces the others a little bit closer together (~107.3°)

NH3

Tetrahedral vs. Trigonal pyramidal

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Tetrahedral geometry around the central atom

Tetrahedral Molecular Shape

Tetrahedral geometry around the central atom

Trigonal Pyramidal Molecular Shape

Tetrahedral vs. Trigonal pyramidal

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On the right, the 4th lone pair, is not seen as part of the actual molecule, yet affects shape.

If another one of the bonding pairs on “trigonal pyramidal” were a lone pair, what is the result?

4 domains on central atom, con’t

BENT

• 4 domains• 2 bonding pairs

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• 2 bonding pairs• 2 lone pairs

• The bonds are forced together still closer (104.5°) by the 2 thick unshared pairs.

H2O

Comparing the 2 “bents”…

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Both bent molecules are affected by unshared pairs –1 pair on the left, 2 on the right.

Other Molecular Geometry

Just for fun, let’s look at some others that we will not study in detail in this course…


detail in this course…

Note that if there are more than five domains Note that if there are more than five domains around the central atom, it must be an exception to around the central atom, it must be an exception to the octet rule!the octet rule!


TRIGONAL BIPYRAMIDAL

• 5 shared pairs

• Three pairs are found in

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• Three pairs are found in one plane (“equator”) 120°apart; the other two pairs are at the “poles,” 180°apart, 90° from the “equator.”

PCl5


SEE-SAW

• 4 shared pairs & 1 unshared pair

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• One of the equator pairs is unshared & pushes the other 2 together.

• The 2 poles are pushed slightly together.

SF4


T-SHAPED

• 3 shared & 2 unshared pairs

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pairs

• 2 of the 3 equator pairs are unshared.

• All 3 remaining pairs are pushed together.

ClF3


LINEAR


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pairs

• All 3 equator pairs are unshared. The 2 remaining pairs are forced to the poles.

XeF2

5 e5 e-- pairs on central atompairs on central atom

5 shared, 0 unshared5 shared, 0 unshared 4 shared, 1 unshared4 shared, 1 unshared

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5 shared, 0 unshared5 shared, 0 unshared



OCTAHEDRAL

• 6 shared pairs

• Each pair repels the others

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• Each pair repels the others equally.

• All angles = 90°

Now, if one of these Now, if one of these pairs was unshared …pairs was unshared …

SF6


SQUARE PYRAMIDAL

• 5 shared pairs & 1 unshared pair

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• 4 shared pairs in one plane; the 5th pair at the pyramid’s top.

If the pair at the top was If the pair at the top was unshared …unshared … IF5


SQUARE PLANAR


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pairs

• The 4 shared pairs are in the same plane; the 2 unshared pairs are 90°from them.

XeF4

6 e6 e-- pairs on central atompairs on central atom


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4 shared, 2 unshared4 shared, 2 unshared

The Valence-Bond Concept•builds on the Lewis proposal that bonding results from electron pairing between neighboring atoms � put into a quantum mechanical context � the results (the valence-bondconcept) were refined by Linus Pauling.

•- used much less now, only organic chemistry and transition metal compounds

The principles

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1. A covalent bond results from the pairing of unpaired electrons in neighboring atoms.

2. The spins of the paired electrons must be antiparallel (one up and one down).

3. To provide enough unpaired electrons in each atom for the maximum bond formation, electrons can be excited to fill empty orbitals during bond formation.

4. The shape of the molecule results from the directions in which the orbitals of the central atom point.

Orbital HybridizationOrbital HybridizationOrbital HybridizationOrbital Hybridization

Solving the Schrodinger wave equation using the valence-bond method requires that the original atomic orbitals be combined into hybrid orbitals that better fit the known geometries.

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The orbital hybridization concept asserts that the wave functions of electrons in atomic orbitals of an atom (usually the central atom of a molecule) can mix together during bond formation to occupy hybrid atomic orbitals.

electrons in these hybrid orbitals are still the property

orbital hybridization.orbital hybridization.orbital hybridization.orbital hybridization.

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of the donor atom.

The number of hybrid orbitals formed will equal the sum of the number of atomic orbitals that are involved in the mixing of wave functions

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Exp 1: BFExp 1: BFExp 1: BFExp 1: BF3333

B =

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matches our experimental findings : 120° angles ; trigonal planar geometry

Exp 2: COExp 2: COExp 2: COExp 2: CO2222

C =

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can be used to explain the linear nature of the CO2 molecule and the presence of two carbon-oxygen double bonds.

Limitation of the Hybridization ConceptLimitation of the Hybridization ConceptLimitation of the Hybridization ConceptLimitation of the Hybridization Concept

-hybridization is simply a mathematical manipulation of wave functions, and we have no evidence that it actually happens.

the hybridization concept is not a predictive tool:

The construction of a set of molecular orbitals, in which the electrons are considered the property of the molecule as a whole, is predictive.

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whole, is predictive.

complexity of the calculations that are needed to deduce molecular shape.

Ex 1: diamond, Ex 1: diamond, Ex 1: diamond, Ex 1: diamond,

Network Covalent Substances

All the atoms are held together by covalent bonds throughout a substanceThe whole crystal is one giant molecule.

Network covalent substances are rare.

sublimes at 4000°C,

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each and every C is bonded in a tetrahedral arrangement to all its neighbors

sublimes at 4000°C,

Ex 2: QuartzEx 2: QuartzEx 2: QuartzEx 2: Quartz

each silicon atom is surrounded bya tetrahedron of oxygen atoms, and each

melts at2000°C.

Ex 2: QuartzEx 2: QuartzEx 2: QuartzEx 2: Quartz

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a tetrahedron of oxygen atoms, and each oxygen atom is bonded to two siliconatoms.

Due to this network covalent bond : 1. m.P >>> 2. extremely hard: 3. insoluble in all solvents.

Amorphous SiliconAmorphous SiliconAmorphous SiliconAmorphous Silicon

Amorphous solids = materials in which the atoms are not arranged in a systematic repeating manner.

Efficiency a-Si = 40 x c-Si at absorbing solar radiation

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c-Sia-Si

Intermolecular Forces

Almost all covalently bonded substances consist of independent molecular units.

only intramolecular forces (the covalent bonds) � no attractions between neighboring molecules � gases !

forces between molecules

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forces between molecules

1. induced dipole attractions/ dispersion forces/London forces

2. Dipoledipole, 3. ion-dipole, 4. hydrogen bonding

only occur in specific circumstances

operates between all molecules

Dispersion (London) ForcesDispersion (London) ForcesDispersion (London) ForcesDispersion (London) Forces

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instantaneous dipoles are created and generate dispersion forces. Instantaneous dipoles forming in one molecule will

generate dipoles in neighbouring molecules due to electrostatic attraction and repulsion. The result of the

development of so many temporary dipoles is a brief, weak force of attraction (dispersion forces)

ElectronegativityElectronegativityElectronegativityElectronegativity

= the power of an atom in a molecule to attract shared electrons to itself.

• reflects the comparative Zeff of the two -atoms on the shared electrons

Relative concept, not a measurable function

there will be a permanentdipole in the molecule.

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measurable function dipole in the molecule.

DipoleDipoleDipoleDipole----Dipole ForcesDipole ForcesDipole ForcesDipole Forces

dipole-dipole attractions effect < induced dipole effect.

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Hydrogen BondingHydrogen BondingHydrogen BondingHydrogen Bonding•the strongest intermolecular force; ≈ 5 to 20 % the strength of a covalent bond.

•hydrogen bond strength : H—F > H—O > H—N

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Molecular Symmetry

Symmetry � to interpret the vibrational spectra of the compounds

Symmetry OperationsSymmetry OperationsSymmetry OperationsSymmetry Operations

- a procedure performed on a molecule that leaves it in a conformation indistinguishable from, and superimposable on, the original conformation

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There are five symmetry operations:1. Identity1. Identity1. Identity1. Identity2. Proper rotation: rotation about an 2. Proper rotation: rotation about an 2. Proper rotation: rotation about an 2. Proper rotation: rotation about an nnnn----fold axis of symmetryfold axis of symmetryfold axis of symmetryfold axis of symmetry3. Reflection through a plane of symmetry3. Reflection through a plane of symmetry3. Reflection through a plane of symmetry3. Reflection through a plane of symmetry4. Inversion through a center of symmetry4. Inversion through a center of symmetry4. Inversion through a center of symmetry4. Inversion through a center of symmetry5. Improper rotation: rotation about an axis followed by a refl 5. Improper rotation: rotation about an axis followed by a refl 5. Improper rotation: rotation about an axis followed by a refl 5. Improper rotation: rotation about an axis followed by a refl ection (real or ection (real or ection (real or ection (real or imagined) perpendicular to that axis

IdentityIdentityIdentityIdentity

The identity operator, E, leaves the molecule unchanged.

Thus, all molecules possess E.

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Proper RotationProper RotationProper RotationProper Rotation

The rotation operation, symbol Cnx, involves rotating the molecule by 360/ndegrees about an axis, symbol Cn, through the molecule.

The value of n represents the number of times the molecule can be rotated during a complete

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the molecule can be rotated during a complete360° rotation while matching the original conformation after each rotation.

Symmetry and Vibrational Spectroscopy

plays an important role in molecular behavior

Ex : in transition metal compounds, molecular symmetry is key in determining the numberand energies of electronic excited states and the probability of an electron being excited into each of those states. These electronic excitations result in the color and color intensity of the compound.

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Transition Metal Complexes

CHAPTER 19


170“Read in the name your God who created “ )I( Inorganic Chemistry 3 )I( Miftahul Khair, M.Sc ([email protected])

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Preface

special interest due to :

• the transition metal compounds come in every‘ color of the rainbow,

Ex : CrCl3.6H2O can be purple, pale green, and dark green

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Ex : CrCl3.6H2O can be purple, pale green, and dark green

Alfred Werner (1893) : transition metal compounds consisted of the metal ions surrounded by the other ions and molecules.

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Transition metals

• is an element that has at least one simple ion with an incomplete outer set of d

electrons.

• Exclude:

1. Group 12 elements-(Zn, Cd, and Hg) -because these metals d10 electron configuration.

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configuration.

2. The Group 3 elements-(Sc, Y, and La) because they almost always exhibit the +3 oxidation state (do electron configuration).

3. postactinoid metals

Introduction to Transition Metal Complexes

A grouping consisted of Transition metal ion covalently bonded to other ions or molecules.

• the diversity of the metal complexes provides the wealth of

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• the diversity of the metal complexes provides the wealth of transition metal chemistry

• Warner : metal ion has combining power; specific number (coordination number) of molecules or ions (ligands) with which transition metals could combine.

Ex: Pt(II), NH3, Cl-, and K+.

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Measurement of electrical conductivity of their solutions and by gravimetric analysis using silver nitrate solution

Stereochemistries

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Tetrahedral : common in Period 4 transition metalssquare planar : Periods 5 and 6.

Ligands

1 coordination site; monodentate

2 coordination site; bidentate,

: the atoms, molecules or ions attached to the metal ion

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etc

•chemdraw

All ligands that form more than one attachment to a metal ion are called chelating ligands

•wide range of oxidation states dependent on the ligands

Nature the ligand Ligands Examples

Stabilize Low Oxidation States

CO, CN- Fe(CO)5

Stabilize "Normal" water, ammonia, [Fe(OH2)6]2+ and

Ligands and Oxidation States of Transition MetalsLigands and Oxidation States of Transition MetalsLigands and Oxidation States of Transition MetalsLigands and Oxidation States of Transition Metals

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Stabilize "Normal" Oxidation States

water, ammonia, and halide ions

[Fe(OH2)6]2+ and

[Fe(OH2)6]3+ .

Stabilize High Oxidation States

fluoride and oxide ions.

[CoF6]2- , [FeO4]

2-,

Isomerism in Transition Metal Complexes

isomer types

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Structural isomers: the bonds to the metal are differentStereoisomer : the bonds to the metal are identical

Structural IsomerismStructural IsomerismStructural IsomerismStructural Isomerism

1. Linkage isomerism1. Linkage isomerism1. Linkage isomerism1. Linkage isomerismligands can form bonds through more than one atom

• E.g : Co(NH3)5Cl2(NO2),

� [Co(ONO)(NH3)5]2+,

� [Co(NO2)(NH3)5]2+,

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2. Ionization isomerism.2. Ionization isomerism.2. Ionization isomerism.2. Ionization isomerism.give different ions when dissolvedin solution

� [Co(NO2)(NH3)5] ,

3. Hydration isomerism.3. Hydration isomerism.3. Hydration isomerism.3. Hydration isomerism. Ag+ precipitation of Cl-test

4. Coordination isomerism.4. Coordination isomerism.4. Coordination isomerism.4. Coordination isomerism.

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4. Coordination isomerism.4. Coordination isomerism.4. Coordination isomerism.4. Coordination isomerism.

both the cation and the anion are complex ions. Ligands interchange betweenthe cation and anion

Stereoisomerism

•Geometric Isomerism

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Optical isomers

• Pairs of compounds in which one isomer is a nonsuperimposable mirror image of the other.

• Found when a metal is surrounded by

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• Found when a metal is surrounded by three bidentate ligands:1,2 diaminoethane : [M(en)3]

n+

• Labile : rapid ligand exchange

• Inert : slow ligand exchange

Naming Transition Metal Complexes

1. Nonionic species are written as one word; ionic species are written as two words with the cation first.

2. The central metal atom is identified by name, which is followed bythe formal oxidation number in Roman numerals in parenthesis,such as( IV) for a +4 state and (-II) for a -2 state.

If the complex is an anion, the ending -ate adds to the metal name or replaces any-ium - en, or –es ending.

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If the complex is an anion, the ending -ate adds to the metal name or replaces any-ium - en, or –es ending.

Ex: cobaltate and nickelate, but Chromate and tungstate (not chromiumate or tungstenate).

For a few metals, the anion name is derived from the old Latin name ofthe element : Ex: ferrate( iron), argentate (silver) ,

cuprate ( copper), and Aurate (gold).

3. The ligands are written as prefixes of the metal name. Neutral ligands are given the same name as the parent molecule, Negative ligands are given the ending -o instead of -e.

Ex: sulfate � sulfato nitrite � nitrito.

Anions with -ide endings completely replaced by -o.Ex: chloride � chloro;

iodide � iodocyanide � cyano;

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cyanide � cyano;hydroxide�hydroxo.

There are three special names : Coordinated water is Commonly named aquaammonia� ammine carbonmonoxide � carbonyl. .

4. Ligands are always placed in alphabetical order (in chemical formulas,

the symbols of anionic ligands always precede those of neutral

ligands).

5. For multiple ligands, the prefixes di-, tri-, tetra- penta-, and hexa- are

Used for two, three, four, five, and six, respectively.

6. For multiple ligands already containing numerical prefixes (such as

1,2-diaminoethane) the prefixes used are bis',tris', and, tetrakis for

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1,2-diaminoethane) the prefixes used are bis',tris', and, tetrakis for

two, three, and four. This is not a rigid rule. Many chemists use

These prefixes for all polysyllabic ligands.

[Pt(NH3)4]Cl2

Tetraammineplatinum(II)chloride.

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cis- diamminedichloroplatinum(II) andtrans-diamminedichloroplatinum(II).

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potassium tetrachloroplatinate(II).

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tris(1,2-

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tris(1,2-diaminoethane)cobalt(III)chloride.

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An Overview of Bonding Theories of TransitionMetal Compounds

Had to account for

1. colors found among the compounds,

2. the wide range of stereochemistries

3. paramagnetism

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3. paramagnetism

The 18-Electron Rule

• octet rule is used to predict the formulas of covalent compounds � occupancy of s and p orbitals

• Similar concept for 18 e- rule. � occupancy of s, p, and dorbitals

• Limitation : valid for very low oxidation state metal

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• Limitation : valid for very low oxidation state metal complexes

• Ex: tetracarbonylnickel(0)Ni(CO)4

Valence-bond theory

• Advantage: can be used for the many non-18 electron complex

• VBT:

the interaction between the metal ion and its ligands to be that of a Lewis acid with Lewis bases,

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that of a Lewis acid with Lewis bases,

but in this case,

the donated ligand electron pairs are considered to occupy the empty higher orbitals of the metal ion.

• can be used for the many non-18 electrons.

[NiCl4]2-

The electron distribution of the free nickel(II) ion

The hybridization and occupancy the higher energy orbitals by electron pairs (open half-headed arrows) of the

Ni= [Ar]3d8

tetrahedral

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1. we can only construct the orbital diagrams once we know from a crystal structure determination and magnetic measurement what the ion shape and number of unpaired electrons actually are.

2. it doesn't explain why the electron pairs occupy higher orbitals, even though there are vacancies in the 3d orbital

3. theory fails to account or the color of the transition metal complexes.

half-headed arrows) of the chloride ligands.However:

Crystal Field Theory

• the transition metal ion is free and gaseous

• and the ligands behave like point charges; and

• there are no interactions between metal orbitals and ligand orbitals.

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orbitals.

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approacho f thc ligandc lectron

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Approaching ligand electrons

orbitalsoriented along the bonding directions will increase in energyand those between the bonding directions will decrease in energy

• seconds step-the loss of degeneracy of the d orbitals�color and magnetic properties explanation.

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Octahedral Complexes

the six ligands are located long

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the six ligands are located long the Cartesian axes

these negative charges along the Cartesian axes the energy dx2-y2

and dz2 orbital will be higher than dxy dxz and dyz.

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The electrons will fall intothe lower energy. This net energy 'decreases known asthe crystal field stabilization energy, or CFSE.

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The energy level splitting depends on four factors:

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