Chapter 6 Jan13

*CHAPTER 6

ELECTRONIC STRUCTURE OF ATOMS

*CONTENT6.1 The wave nature of light6.2 Quantized Energy and Photons6.3 Bohrs Model of the Hydrogen Atom6.4 The Wave Behaviour of Matter6.5 Quantum Mechanics and Atomic Orbitals6.6 Representations of Orbitals6.7 Orbital in Many-Electron Atoms6.8 Electron Configurations6.9 Electron Configurations and The Periodic Table

*Learning Outcomesshould be able to name the orbital of an electron from the n, l, ml values givenAble to write the electronic configuration of any elementAble to calculate the energy associated with electron movement in atom (Bohr model)

*INTRODUCTIONIn the Periodic Table, elements that exhibit similar properties are placed together in the same column. E.g.i. Na and K (Group 1A) are soft reactive metals.ii. He and Ne (Group 8A) are unreactive gases.

The fundamental reason for this similarity lies on the behaviour of electrons in atoms.

*Cont: INTRODUCTIONWhen atoms react it is the electrons that interact (share/transfer electron).

Therefore, we must understand the behaviour of electrons.

Arrangement of electrons in atom is called electronic structure (configuration).

*Cont: INTRODUCTIONELECTRONIC STRUCTURE

Electronic structure refers to:

- No. of electrons that an atom possesses.- Electron distribution around the atom.- Electrons energies.

*6.1 The Wave Nature of Light

When you sense the warmth of a fire, you are feeling what scientists call radiation or electromagnetic radiation.

Electromagnetic radiation carries energy through space and is therefore also known as radiant energy.

*Cont: 6.1 The Wave Nature of LightFire gives off light (visible radiation) and heat (thermal radiation).

Both types of electromagnetic radiation exist in the form of electromagnetic waves.

All waves have wavelength, (lambda) and amplitude, A and frequency, (nu) characteristics.

*Cont: 6.1 The Wave Nature of LightFrequencyThe frequency (nu) of a wave is the number of cycles that pass through a given point in one second.

Unit of frequency is Hertz (Hz).

1 Hz = 1 cycle per second or (1 s-1).

*Cont: 6.1 The Wave Nature of LightWavelengthThe wavelength, (lambda) is the distance between successive peaks (or valleys)

Unit of depends on the type of radiation.

*Cont: 6.1 The Wave Nature of Light Common units for Electromagnetic Radiation: unit symbol length(m) type of radiationAngstrom 10-10x raynanometer nm 10-9 UV, visiblemicrometer m 10-6 IRmilimeter mm 10-3 IRcentimeter cm 10-2 microwavemeter m 1 TV, radio

*Cont: 6.1 The Wave Nature of LightRelationship between and : = c / or c = where c is the speed of light, 3.00 108 ms-1.

*Example 1Calculate the wavelength of an electromagnetic radiation that has a frequency of 9.22 1017 s-1.

c = ; = c / = 3.00 108 ms-1 9.22 1017 s-1 = 3.25 10-10 m

*6.2 Quantized Energy and Photons3 Phenomena involving matter and light were confounding to physicists of the early 20th century:

a) Blackbody radiationb) The photoelectric effectc) Line spectra

*6.2.1 Blackbody RadiationWhen solids are heated, they emit radiation.

Eg:1000 K, heat a coal soft red glow1500 K, light is brighter and more orange 2000 K, light brighter and whiter

*Cont: 6.2.1 Blackbody RadiationMax Planck (1900) proposed that the minimum amount of radiant energy (E) that an object can release or absorb is related to the frequency ( ) of the radiation.E=h where h = Plancks constant

The minimum (smallest quantity) energy that can be emitted or absorbed as electromagnetic radiation is called a quantum of energy.

*Cont: 6.2.1 Blackbody Radiation

For a single quantum: E = h h = Plancks constant = 6.63 10-34 Joule-seconds (Js) = frequency

*Cont: 6.2.1 Blackbody RadiationPlancks quantum theory

energy is quantised (discreet values)

always emitted or absorbed in whole-number multiples of h, e.g. h, 2h, 3h...

*Example 2Calculate the smallest increment of energy, the quantum of energy, that an object can absorb from yellow light whose wavelength is 589 nm. = c/ = 3.00 108 ms-1 589 10-9 m = 5.09 1014 s-1E= h= (6.63 10-34 Js) (5.09 1014 s-1)= 3.37 10-19 J

*Cont: 6.2.1 Blackbody RadiationPlancks theory tells us that an atom or molecule emitting or absorbing radiation whose wavelength is 589 nm cannot lose or gain energy by radiation except in multiples of 3.37 10-19 J. (1 h = 3.37 10-19 J)If one quantum of radiant energy supplies 3.37 10-19 J, then one mole of these quanta will supply: (6.022 1023 quanta) (3.37 10-19 J / quantum ) = 2.03 105 J.

*6.2.2 The Photoelectric EffectAlbert Einstein used the quantum theory to explain the photoelectric effect.

Photoelectric effect - light shinning on a clean metal surface causes the surface to emit electrons.

Einstein proposed that a beam of light consists of a collection of small particles of energy, called photons.

*Cont: 6.2.2 The Photoelectric EffectEach photon must have an energy proportional to the frequency of the light thus E = h.

When a photon strikes the metal, its energy is transferred to an electron in the metal.

The electrons will only be ejected once the threshold frequency is reached.

*Cont: 6.2.2 The Photoelectric Effect

An electron is ejected if it gains sufficient energy (energy threshold) to overcome the forces binding the electron (attractive forces that hold it within the metal).

If a photon has more than the minimum energy required to free an electron, the excess appears as the kinetic energy of the emitted electron.

QUESTIONIt requires 183.7 kJ/mol to eject electrons from cesium metal. What is the minimum frequency of light necessary to emit electrons from cesium via photoelectric effect?What is the wavelength of this light?If cesium is irradiated with light of 416 nm, what is the maximum possible kinetic energy of the emitted electrons?

*

*6.3 Bohrs Model of The Hydrogen Atom6.3.1 Line Spectra

Radiation composed of one wavelength is called monochromatic.

Most radiation sources consist of many different wavelengths.

Dispersion of radiation into its component wavelengths produces a spectrum.

*Cont: 6.3.1 Line Spectra

A spectrum from white light can be separated into a continuous range of colours (rainbow) covering all wavelengths called continuous spectrum.

*Cont: 6.3.1 Line Spectra

When different gases are placed in a tube, under reduced pressure and high voltage, the gases emit different colours of light

When these lights are passed through a prism, discrete lines are observed- line spectrum.

H2 (purple), Ne (red-orange), Na vapor (yellow), N2 (orange), Hg vapor (blue)

*6.3.2 Bohrs ModelNiels Bohr proposed a model for the hydrogen atom that enabled scientists to explain hydrogen line spectrum.

Bohr put forward 3 postulates:An electron in an atom is allowed certain energies known as stationary states (orbits). Radiation and absorption of photons is due to changes in stationary states.An electron moves in circular orbits (allowed energy) around the nucleus and does not spiral into the nucleus.

*Cont: 6.3.2 Bohrs ModelIn this model, the energy of the hydrogen atom, E, depends on the value of a number, n, called the principal quantum number { n is a positive integer: n = 1,2,3........}En = (-hcRH) RH = Rydberg constant = 1.10 107 m-1 h = Plancks constant, c = speed of lightTherefore: En = (-2.18x10-18J) The energies of the electron of a hydrogen atom are negative for all values of n.

*Cont: 6.3.2 Bohrs ModelThe lower (more negative) the energy is, the more stable the atom will be.

The energy is lowest (most negative) for n = 1; ground state.

Other values of n correspond to excited states of the atom.

*Cont: 6.3.2 Bohrs ModelRadius of the stationary states (orbits) increases as n increases and is proportion to n2.

The radius of the orbit n=2 is 22 =4 times larger than the radius of the orbit for n=1.

When n=, (the highest energy state), E=0 indicating that electron is removed from the nucleus.

*Cont: 6.3.2 Bohrs ModelWhen an electron moves from a lower (more negative) energy state to a higher (less negative) energy state, light is absorbed.

When an electron moves from a higher energy state to a lower energy state, light is emitted.

Bohrs model is applicable only for atoms or ions with a single electron (H, He+ and Li2+)

*Cont: 6.3.2 Bohrs ModelE = Ef - Ei where Ef = final state energy Ei = initial state energy

where ni = initial state quantum no. nf = final state quantum no.(ni > nf , energy is emitted, has a negative value)(nf > ni , energy is absorbed, has a positive value)

*Cont: 6.3.2 Bohrs ModelCalculation (n = 2 to n = 3)

*Cont: 6.3.2 Bohrs ModelCont: Calculation (n = 2 to n = 3)

One of the wavelengths observed in hydrogen line spectrum.

*6.3.3 Summary of Atomic Modeli.First orbit has n = 1 (closest to the nucleus).ii.The furthest orbit has n = and corresponds to E = 0iii. Electrons can move between orbits by absorbing or emitting energy in quanta (h, 2h, 3h)iv. Emission or absorption of radiant energy from a hydrogen atom occurs when an electron moves from one orbit to another.

*Cont: 6.3.3 Summary of Atomic Modelv. If the electron transfer from an initial state with energy Ei to a final state with energy Ef:E = Ef - Ei where E = h = -hcRH(1/n 2)

vi. When nf ni , energy is absorbed, is positive.vii. When nf ni , energy is emitted, is negative.

*Cont: 6.3.3 Summary of Atomic Model

viii. If a calculated frequency is 2.5 1015 s-1, that means that light of frequency 2.5 1015 s-1 is absorbed during the transition.

ix. If a calculated frequency is -2.5 1015 s-1, that means that light of frequency 2.5 1015 s-1 is emitted during the transition.

*Example 3

What are the postulates of Niels Bohrs model for the hydrogen atom?

*Example 3 (Answer)The postulates are:An electron moves in a circular path about the nucleus and does not spiral into the nucleus. The energy of an electron can have only certain allowed values; its energy is quantized.An electron can move from one circular path to another circular path only when it absorbs or emits a photon.

*Example 4

1.What is the meaning of the negative sign in the equation En = -hcRH(1/n2)?

2.What is an excited state?

3.When an atom is ionised, what is the final principal quantum number of the electron?

*Example 4 (Answer)Zero point of the atoms energy is when the electron is completely removed from the nucleus. Thus E = 0 when n = , so E 0 for any n value smaller than .Any orbit of the hydrogen atom with n 2 is considered a higher energy orbit and is said to be an excited state.Ionisation corresponds to a transition to a final state of n =

*6.4 The Wave Behaviour of Matter

Radiation appears to have either a wavelike or particlelike (photon) character.

If radiant energy could, under appropriate conditions, behave as though it were a stream of particles, could matter, under appropriate conditions, possibly show the properties of a wave?

*Cont: 6.4 The Wave Behaviour of Matter

Using Einsteins and Plancks equations, Louis de Broglie derived: = h h = Plancks constant mv v = velocity m = mass He proposed that the characteristic wavelength of the electron or of any other particle depends on its mass, m, and velocity, v; where the quantity mv is called momentum.

*6.4.1 Heisenberg Uncertainty Principle

This concept of waves and particles applies to low mass, high speed objects.

Discovery of the wave properties of the electron led to Heisenbergs uncertainty principle, which indicates that the position and momentum of an electron cannot both be known with absolute certainty simultaneously.

Thus it is not appropriate to imagine the electrons as moving in well-defined circular orbits about the nucleus.

*6.5 Quantum Mechanics and Atomic OrbitalsErwin Schrdinger proposed the wavefunction equation , (psi) which contains both wave and particle terms.

The location of the electron cannot be determined exactly but the probability of it being at a particular position is given by the probability density, 2.

His work opened a new way of dealing with subatomic particles known as:Quantum mechanics or wave mechanics.

*Cont: 6.5 Quantum Mechanics and Atomic OrbitalsNote:

Bohrs model - the electron is in a circular orbit of same particular radius about the nucleus.

In quantum mechanics - the electrons exact location cannot be determined but the probability of finding an electron at a given point in space can be calculated.

*6.5.1 Orbitals and Quantum NumbersSolving the Schrdingers equation will give us wave functions and energies for the wave functions.

These allowed wave functions are called orbitals.

Each orbital describes a specific distribution of electron density in space.

Light + interaction with matter (blackbody radiation, photoelectric effect, line spectra)

*


Quantum theory(Bohrs model, wave behaviour of matter, Heisenberg Uncertainty Principle)

*



Quantum Mechanics (involve math, calculations, equations)

*



Quantum Mechanics (involve math, calc, equations)

QUANTUM NUMBER(tell us how to write electron configuration)

*

*Cont: 6.5.1 Orbitals and Quantum NumbersIn Quantum Mechanics, the motion of an electron about its nucleus and its energy is characterised by 4 quantum numbers to describe an orbital.

n- principal quantum numberl- second quantum number/azimuthal quantum numberml - magnetic quantum numberms electron spin quantum number

*1. The principal Quantum Number, n n = 1, 2, 3.. (positive integer)

Represents the main energy levels for the electron as a shell in space where the probability of finding an electron is high.

As n increases:- the orbital becomes larger- the electron is further away from the nucleus- the electron has higher energy (less tightly bound to the nucleus)

*2. The Second Quantum Number or the Azimuthal Quantum Number, l

l = (0, 1, 2, 3,......n -1) for each value of n.

Defines the shape of orbital.

Specifies subenergy levels within the main energy levels called subshells.

Subshell s, p, d and f are designated by the l values of 0, 1, 2 and 3, respectively.

*3. The Magnetic Quantum Number, mlDescribe the orientation of a single atomic orbital in space.The number of different permissible orientations of an orbital depends on the l value of a particular orbital.The ml quantum number has permissible values of -l to +l, including zero.

*Cont: 3. The Magnetic Quantum Number, mlEg. when l = 0, only one allowed value for ml which is 0 (zero).

when l = 1, three permissible values for ml which are -1, 0, +1

For a given value of l, there are (2l + 1) values of ml.

*4. The Electron Spin Quantum Number, msIt was postulated that electrons have an intrinsic property called electron spin, given by the electron spin quantum number, ms.

The values: - , + (two allowed spin directions)

The spin quantum number has minor effect on the energy of an electron.

*Cont: 4. The Electron Spin Quantum Number, msPauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms)

This indicates that each orbital can only be occupied the most by two electrons of opposite spin.

*6.5.2 Summary of Quantum NumbersThe collection of orbitals with the same value of n is called an electron shell.

Eg: All orbitals that have n = 3 are said to be in the third shell.

The set of orbitals that have the same n and l values is called a subshell.

*Cont: 6.5.2 Summary of Quantum NumbersThe shell n consists of exactly n subshells.

Each subshell has value l from 0 to n-1.

First shell n =1; consists of one subshell 1s (l = 0)

*Cont: 6.5.2 Summary of Quantum NumbersSecond shell n = 2; consists of two subshells 2s (l = 0) and 2p (l = 1)

Third shell n = 3; consists of three subshells 3s (l = 0), 3p (l = 1), and 3d (l = 2)

Forth shell n = 4; consists of four subshells 4s (l = 0), 4p (l = 1), 4d (l = 2) and 4f (l = 3)

*Cont: 6.5.2 Summary of Quantum NumbersEach subshell consists of a specific number of orbitals.

For l value, there are (2l + 1) allowed values of ml.

Each orbital corresponds to a different allowed value of ml.

*Cont: 6.5.2 Summary of Quantum NumbersEg. Subshell s ( l = 0) consist of one orbital = (2l + 1) = ( 2(l) + 1) = 1Eg. Subshell p (l = 1) consist of three orbitals = (2(l) + 1) = 3 and so forth.

*Cont: 6.5.2 Summary of Quantum Numbers Total no of orbital in a shell is n2.n total no of orbitals in shell1 12 43 94 16

*Example 5Write the shorthand notation (for eg 2s) for the orbitals described by the following quantum numbers:

a) n = 2, l = 1, ml = 0b) n =3, l = 2, ml = 1c) n = 4, l = 3, ml = 2

*Example 5 (Answer)a) n = 2l = 1 corresponds to p orbital. the orbital designation is 2pml = 0 could be 2px or 2py or 2pz

b) n = 3 l = 2 corresponds to d orbital. the orbital designation is 3d. ml = 1 could be one of the 5 orbitals.

*Example 5 (Answer)c) n = 4 l = 3 correspond to f orbital. the orbital designation is 4f. ml = 2 could be one of the 7 orbitals.

*Example 6Why is it not possible for a hydrogen orbital to have the quantum numbers n =3, l = 2, ml = 3 associated with it?

*Example 6Why is it not possible for a hydrogen orbital to have the quantum numbers n =3, l = 2, ml = 3 associated with it? For a given l value only certain ml values are allowed.If n = 3; l = 0, 1, 2 (as l = 0 to n-1) ml = l to -l = 2, 1, 0, -1, -2

Thus ml = 3 is not possible for a hydrogen orbital with an l = 2 value.

*Example 7Which of the following is a correct set of quantum numbers for an electron in a 3d orbital?

A.n = 3, l = 0, ml = -1B.n = 3, l = 1, ml = +3C.n = 3, l = 2, ml = 0D.n = 3, l = 3, ml = +2

*Solution

nlml

30 01 -1, 0, +12 -2, -1, 0, +1, +2

*6.6 Representations of Orbitals6.6.1 The s Orbitals

The s orbital is spherical.As n increases, the s-orbitals get larger.1s orbital is the lowest-energy orbital.The intermediate regions where 2 is zero are called nodes.A node is a region in space where the probability of finding an electron is zero.No. of nodes increases as n increases.

*6.6.2 The p OrbitalThere are three p-orbitals ; px, py and pz.The letters correspond to allowed values of ml of -1, 0 and +1.Each shell beginning with n = 2 has three p orbitals.n = 2 ;l = 0, 1 (s and p orbitals)ml = 0 ; -1, 0, +1 2s 2px 2py 2pzSize of p orbitals increases as n increases.All p orbitals have a node at the nucleus.

*6.6.3 The d and f Orbitals When n 3, we encounter the d orbitals (l = 2)When n 4, we encounter the f orbitals (l = 3)There are 5d and 7f - orbitals.n = 4, l = 0, 1, 2, 3, = s, p, d, f

For d orbital l = 2, ml = -2, -1, 0, 1, 2 (5 orbitals)For f orbital l = 3, ml = -3, -2, -1, 0, 1, 2, 3 (7 orbitals)

*Cont: 6.6.3 The d and f Orbitals

Three of the d-orbitals lie in a plane bisecting the x, y and z axes (dxy, dyz, dxz).

Two of the d-orbitals lie in a plane aligned along the x, y and z axes (dx2-y2, dz2).

f orbitals will not be discussed here.

*6.7 Orbitals in Many-Electron AtomsIn hydrogen, the energy of an orbital depends only on its principal quantum number, n (for n=3; 3s, 3p and 3d subshells have the same energy).

In many-electron atom, the electron-electron repulsions cause different subshells to be at different energies although they are from the same principal quantum number.

*6.7.1 Effective Nuclear Charge (Screening Effect)In many-electron atom, each electron is simultaneously:i. Attracted to the nucleus (positively charged).ii. Repelled by other electrons.

We can estimate individual electron energy by considering how it interacts with the average environment created by the nucleus and other electrons.

*Cont: 6.7.1 Effective Nuclear Charge (Screening Effect)Any electron density between the nucleus and the electron of interest will reduce the nuclear charge acting on that electron.

Net positive charge attracting the electron is called effective nuclear charge, Zeff. Zeff = Z - S Z = number of protons S = average number of electrons

*Cont: 6.7.1 Effective Nuclear Charge (Screening Effect)The positive charge experienced by outer-shell electrons is always less than the full nuclear charge.

Reasons:i. The inner-shell electrons partly offset the positive charge of the nucleusii. The inner electrons shield the outer electron from the full charge of the nucleus :screening effect (or shielding effect)

*6.7.2 Energies of OrbitalsAs we move outward from the nucleus, the screening effect depends on its electron distribution.

For a given value of n, electron distribution differs for each subshell.

*Cont: 6.7.2 Energies of OrbitalsEg : n = 3 (3s, 3p, 3d)- The closeness of the electrons to the nucleus in the orbitals decrease:3s > 3p> 3d

- As a result, the shielding effect experienced by the electrons increase:3s < 3p< 3d

- Therefore, the Zeff experienced by the electrons in the orbitals decrease:3s > 3p> 3d

*Cont: 6.7.2 Energies of OrbitalsGenerally, in a many-electron atom, for a given value of n, Zeff decreases with increasing value of l.The energy of an electron depends on Zeff.All orbitals of a given subshell, such as 3d orbitals have the same energy.Orbitals with the same energy are said to be degenerate.

*6.8 Electron ConfigurationThe way in which the electrons are distributed or arranged among orbitals of various energy levels around an atomic nucleus. The most stable or ground state of the electron configuration : the electrons in the lowest possible energy states.The most stable ground state is 1s orbital.Pauli exclusion principal : there can only be a maximum of two electrons in any single orbital.

*Cont: 6.8 Electron ConfigurationThe orbitals are filled in the order of increasing energy - 2 electrons per orbital.The maximum number of electrons which can accommodate in each shell in an atom is defined by different sets of the four quantum numbers and is 2n 2.

Eg: There can only be maximum of 2 electrons for the first principal shell, n =1 (n=2, maximum electrons=8; n=3, maximum electrons=18).

*Cont: 6.8 Electron ConfigurationElectron configurations are written that lists the principal quantum number first, followed by an orbital letter s, p, d or f.The order in which the electrons fill up the orbitals is as follows:1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s2 5f 14 6d 10 7p 6A convenient way to memorise this order is by following the arrows from tail to head.

*Cont: 6.8 Electron ConfigurationWe can also show the arrangement of the electrons as:Hydrogen 1s 1 -

Helium 1s 2 -

Lithium 1s 2 2s 1 -

This is called an orbital diagram.

*Cont: 6.8 Electron ConfigurationOrbital diagram.Each orbital is represented by a box.Each electron is represented by a half arrow.

ms = + ms = - (spin-up)(spin-down)

*6.8.1 Periods 1, 2 and 3How the electron configurations of the elements change as we move from element to element across the Periodic Table. All the three 2p orbitals are of equal energy.For carbon, the 6th electron must go into 2p orbital.

*Cont: 6.8.1 Periods 1, 2 and 3Follow Hunds Rule - for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximised.This means that the orbitals of a subshell must be occupied singly and with parallel spins before they can be occupied in pairs.

*Cont: 6.8.1 Periods 1, 2 and 3For Nitrogen we have 7 electrons.The orbital diagram will be as follows:

-The electrons will occupy orbitals singly to the maximum extent possible to achieve its lowest energy - minimising electron-electron repulsion.

*Cont: 6.8.1 Periods 1, 2 and 3As for sodium (Z = 11), it has a single 3s electron beyond the stable configuration of neon.

Abbreviation of the electron configuration of sodium is as follows:Na [Ne] 3s 1

*Cont: 6.8.1 Periods 1, 2 and 3Symbol [Ne] represents the electron configuration of the 10 electrons of Neon 1s 2 2s 2 2p 6.This system helps to focus on the outermost electrons of the atom.The outer electrons are the ones largely responsible for the chemical behaviour of an element.

*Cont: 6.8.1 Periods 1, 2 and 3Eg : Li [He] 2s1

Li and Na have the same type of outer-shell electron configuration.The outer-shell electrons are often referred to as valence electrons.The electrons in the inner shells are the core electrons.All the members of the alkali metal group (1A) have a single s electron beyond a noble gas configuration.

*6.8.2 Period 4 and Beyond The next element following Argon in the Periodic Table is potassium (K).

Argons electron configuration is 1s2 2s2 2p6 3s2 3p 6.

Potassium is a member of alkali metal group ie K has a single s electron beyond Argon configuration.

This means that the highest-energy electron has not gone into 3d orbital.

*Cont: 6.8.2 Period 4 and Beyond4s orbital is lower in energy than that of 3d.

4s orbital is filled first followed by 3d orbitals.

Eg: 19K - 1s2 2s2 2p6 3s2 3p6 4s1 (19K - [Ar] 4s 1) 20Ca - 1s2 2s2 2p6 3s2 3p6 4s2 (20Ca - [Ar] 4s2)

*Cont: 6.8.2 Period 4 and Beyond 21Sc - 1s2 2s2 2p6 3s2 3p6 4s2 3d 1 (21Sc - [Ar] 4s2 3d 1) 30Zn - [Ar] 4s 2 3d 10

*Sc represents the start of transition metal series

Follow Hunds Rule, electrons are added to the 3d orbitals singly until all 5 orbitals have one electron each.

*Cont: 6.8.2 Period 4 and BeyondAdditional electrons are placed in the 3d orbitals with spin pairing.

Electrons having opposite spins are said to be paired when they are in the same orbital.

Upon completion of 3d transition series, the 4p orbitals begin to be filled.

*6.9 Electron Configurations and The Periodic TableElectron configuration of the elements are related to their location in the Periodic Table.

Elements with the same pattern of valence electron configuration are arranged in the same column.

Eg: All 2A elements have ns2 outer configuration.

All 3A elements have ns2 np1 outer configuration.

*Cont4

*Cont: 6.9 Electron Configurations and The Periodic TableThe total number of orbitals in each shell is equal to n 2 : 1, 4, 9 or 16.The total number of electrons a particular shell can accommodate is 2n2 electrons : 2, 8, 18 or 32.Consider the Periodic Table:- First row has 2 elements.- 2nd and 3rd rows have 8 elements.- 4th and 5th rows have 18 elements.- 6th row has 32 elements (including lanthanides).Therefore, the Periodic Table is the best guide to the order in which orbitals are filled.

*Cont: 6.9 Electron Configurations and The Periodic TableElement can be grouped in terms of the type of orbital into which the electrons are placed.1A and 2A elements (s-block) - the outer-shell s orbitals are being filled.3A to 8A elements (p-block) - the outer most p orbitals are being filled.s-block and p-block contain the representative (or main-group) elements.Transition metals (d-block) - the outer most d orbitals are being filled.

*Cont: ELECTRON CONFIGURATIONS AND THE PERIODIC TABLE4

*SUMMARY AND KEY TERMSElectrons in an atom show two intrinsic spin orientations in the presence of a magnetic field.The two spin orientations are in opposite directions and are quantized.Electron spin is defined by the electron-spin quantum number, ms, which has two possible values, ms = +1/2 , ms = -1/2Pauli exclusion principle states that:No two electrons in an atom can have the same set of four quantum numbers n, l, ml and ms.

*Cont: SUMMARY AND KEY TERMSAn orbital can hold a maximum of two electrons and these electrons must have opposite spins.

Subshells (for example 2s and 2p) do not have the same energies as they do in a hydrogen atom.

Orbitals within a subshell (for example 2px, 2py and 2pz) possess the same energy.

*Cont: SUMMARY AND KEY TERMSElectrons occupy orbitals in order of increasing energy.

All orbitals that are equal in energy (degenerate) are filled with electron first, before the next level of energy begins to fill up.

(A few exceptions exist : Cr and Cu)

*Example 8Write electron configuration for Ti (atomic no.= 22Solution:Step 12He: 1s 210Ne: 1s 2 2s 2 2p 618Ar: 1s 2 2s 2 2p 6 3s 2 3p 636Kr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6

Ar: 18 electronsTi: 22 electrons [Ar] + 4 electrons [Ar] 4s 2 3d 2

*Example 8 (cont)Remember of a few exceptions:

24Cr : [Ar] 4s 1 3d 5

29Cu : [Ar] 4s1 3d10

47Ag : [Kr] 5s1 4d10

*

END of CHAPTER 6

Documents

Chapter 6 Jan13