Chemistry 281(01) Winter 2014

Chapter-1-1Chemistry 281, Winter 2014, LA Tech

CTH 277 10:00-11:15 am

Instructor: Dr. Upali Siriwardane

E-mail: [email protected]

Office: 311 Carson Taylor Hall ; Phone: 318-257-4941;

Office Hours: MTW 8:00 am - 10:00 am;

TR 8:30 - 9:30 am & 1:00-2:00 pm.

January 14, 2014 Test 1 (Chapters 1&,2),

February 6, 2014 Test 2 (Chapters 3 &4)

February 25, 2014, Test 3 (Chapters 5 & 6),

Comprehensive Final Make Up Exam: February 27, 2012

9:30-10:45 AM, CTH 311.

Chemistry 281(01) Winter 2014

mailto:[email protected]


Chapter 1. Atomic SturctureChapter 1. Atomic structure 3

The origin of the elements 3

1.1 The nucleosynthesis of light elements 5

1.2 The nucleosynthesis of heavy elements 6

1.3 The classification of the elements 8

The structures of hydrogenic atoms 10

1.4 Spectroscopic information 10

1.5 Some principles of quantum mechanics 11

1.6 Atomic orbitals 12

Many-electron atoms 18

1.7 Penetration and shielding 18

1.8 The building-up principle 20

1.9 Atomic parameters


Origin of Elements in the Universe Scientists have long based the origin of our Universe on the Big Bang Theory. According

to this theory, our universe was simply an expanding fairly cold entity consisting of only

Hydrogen and Helium during it's incipient stages. Over the expanse of many years, and

through a continuing process of fusion and fission, our universe has come to consist of

numerous chemical elements, four terrestrial planets (Earth, Mars, Venus, and Mercury),

and five giant gas planets (Saturn, Jupiter, Neptune, Pluto, and Uranus).


Eight Steps in the History of the Earth1. The Big Bang2. Star Formation 3. Supernova Explosion4. Solar Nebula Condenses 5. Sun & Planetary Rings Form6. Earth Forms 7. Earth's Core Forms 8. Oceans & Atmosphere Forms


Nuclear Chemistry• Fusion is lighter nuclei coming together to form

heavier.• Fission is heavier nuclei breaking in to lighter

nuclei.• Mass is not conserved E=mc2

• Nuclear reactions are balanced by A (mass) and Z (atomic) number.

• Energy released is E=mc2, m is mass defect in amu mutiplied by the conversion factor (931.5 MeV/amu)

• Binding energy of nuclei expressed in Mev/nucleons


Balancing Nuclear Equations


Nuclear Binding EnergyThe binding energy of a nucleus is a measure of how

tightly its protons and neutrons are held together by the nuclear forces. The binding energy

per nucleon, the energy required to remove one neutron or proton from a nucleus, is a function

of the mass number A. (Dm) –mass defect(Dm) = Mass of Nuclide - mass of (p + n +e ) Proton mass: 1.00728 amuNeutron mass: 1.00867 amu 931.5 MeV/amuElectron mass: 0.00055 amuMass defect (Dm), then multiply by


Bonding Energy Curve


Nuclear Fusion Reactions• Nuclear energy, measured in millions of electron

volts (MeV), is released by the fusion of two light nuclei, as when two heavy hydrogen nuclei, deuterons (2H), combine in the reaction


Nuclear Fission Reactions• Nuclear energy is also released when the fission

(breaking up of ) of a heavy nucleus such as U is induced by the absorption of a neutron as in


Origin of the Elements: Nucleosynthesis•Elements formed in the universe's original stars were made from hydrogen gas condensing due to gravity. These young stars "burned" hydrogen in fusion reactions to produce helium and the hydrogen was depleted. Reactions such as those below built up all the heavier elements up to mass number 56 in the periodic table.•When the stars got old they exploded in a supernova, spreading the new elements into space with high flux of neutrons to produce heavy elements by neutron capture.


Nuclear Burning


Supernova Explosion


The nucleo-synthesis of light elements • Stellar nucleo-synthesis

Elements carbon to Iron is form by nuclear fusion in stars after all H is converted to He.

• Double star SupernovaWhite dwarf (dense ball of carbon/oxygen) steals material from another star and get heated releasing huge energy. It goes to nuclear overload and carbon/oxygen suddenlyfuses to iron and it explodes known as type 1a supernova. Most of the elements up to iron in the universe


The nucleosynthesis of heavy elements Havier elements are formed during Supernova explosion.

Giant one star supernova explosionsA heavier star buns all its H and nuclear burning goes faster and forms layer after layers of new elements with increasing mass number up to iron. Core collapses and become denser and the star explodes. Iron capture neutrons and all heavier elements beyond iron.Corpse of supernova explosion leaves a core neutrons. Rotating neutron produces EM pluses creating a pulsar

Hypernova explosions: g-ray bursts


Cosmic Abundances


Terrestrial Abundances

http://upload.wikimedia.org/wikipedia/en/5/56/Relative_abundance_of_elements.png


Stability of the Elements and Their Isotopes

P/N RatioWhy are elementsWith Z > 82 areUnstable?


Terrestrial Abundances


Magic Numbers• Nuclei with either numbers of protons or

neutrons equal to Z, N =2 (He), 8(O), 20 (Ca), 28(Si), 50(Sn, 82(Pb), or 126(?)(I)

• exhibit certain properties which are analogous to closed shell properties in atoms, including

• anomalously low masses, high natural abundances and high energy first excited states.


The classification of the elements• Dobereiner Triads• Newlands called the Law of Octaves• Mendeleyev’s periodic table• Lothar Meyer’s atomic volume curves• Glen Seaborg atomic number and long form


Dobereiner Triads

Cl 35.5 Li 7 S 32

Br 79 Na 23 Se 79

I 127 K 39 Te 128


Newlands’ Law of octaves

Octaves 1 Li Be B C N O F

Octaves 2 Na Mg Al Si P S Cl


Lothar Mayer’s atomic volume curves


Mendeleyev’s Periodic Table


Long Form of Periodic Table


What is periodic table? Describe its use in chemistry?

All elements in a group have similar chemical properties

Group I- alkali metal:Li, Na, K Rb, Cs, Fr Common ele.n conn: ns1

Group II- alkaline earth metals:Be, Mg, Ca, Sr, Ba, Ra: Common ele.n conn: ns2

Group VII- Halogens: Cl, Br, I, At: Common ele.n conn:ns2 np5

Group VIII- Noble gases:He, Ne, Ar, Kr, Xe, Rn: Common ele.n conn ns2 np6


Chemical properties and the periodic tableElectron configurations help us understand changes in atomic radii, ionization energies, and electron affinities.Various trends in reactivity can be observed.

• Main group metals become more reactive as you go down a group.

• Reactivity of nonmetals decreases as you go down a group.

• Transition metals become less reactive as you go down a group.


Other ways of numbering groups in the periodic table• Several methods are used for numbering periodic table

groups• American chemists preferred method.

• The IUPAC old system.• The IUPAC current system.

• The American Chemical Society (ACS) has also adopted the current IUPAC system.


Other numbering systems

H

Li

Na

K

He

Be

Mg

Ca ZnCuTiSc NiCoFeMnCrV Ga KrBrSeAsGe

Al ArClSPSi

B NeFONC

IA IIA IIIA IIIA IVA VA VIA VIIA 0

IIIB IVB V B VIB VIIB VIII B IB

IIB

1

2

3

4

1 2 13 14 15 16

17 18

IA IIA III B IVB VB VIB VIIB

VIIIB

3 4 5 6 7 8 9 10

11 12

IIIA IVA VA VIA VIIA VIIIA IB IIB

Previous IUPAC

Current IUPAC and ACS

Preferred US


The structures of hydrogenic atoms :Bohr Theory

• The Bohr model is a ‘planetary’ type model.

• Each principal quantum represents a new ‘orbit’ or layer.

• The nucleus is at the center of the model.


Emission Spectrum of Hydrogen• Bohr studied the spectra produced when atoms were excited in a gas discharge tube.

He observed that each element produced its

own set of characteristic lines.


Emission Spectrum of Hydrogen

• Line Spectrum• Energy is absorbed when an electron goes from a

lower(n) to a higher(n) • Energy is emitted when an electron goes from a

higher(n) to a lower(n) level • Energy changed is given by:DE = Ef - Ei

• or DE = -2.178 x 10-18 [1/n2f - 1/n2

i] J• DE is negative for an emission and positive for an

absorption • DE can be converted to l or 1/ l by l = hc/E.


Bohr model of the atom• The Bohr model is a

‘planetary’ type model.

• Each principal quantum represents a new ‘orbit’ or layer.

• The nucleus is at the center of the model.

• RH = 2.178 x 10-18 JEn = -

En = RH


What is Bohr’s Atomic model? • explain emission spectrum of hydrogen atom• applied the idea of Quantization to electrons to orbits• energies of these orbits increase with the distance

from nucleus.• Energy of the electron in orbit n (En):• En = -2.178 x 10-18 J (Z2/n2)• En = -2.178 x 10-18 J 1/n2; Z=1 for H


Bohr model of the atomBalmer later determined an empirical relationship that described the spectral lines for hydrogen.

DE = - 2.178 x 10-18

J ( )1

nf2

1

ni2-

nf = 2 ni = 3,4, 5, . . . Blamer series

Spectra of many other atoms can be described by

similar relationships.

En = -


Paschen, Blamer and Lyman Series


Calculation using the equation: E = -2.178 x 10-18 (1/nf

2 - 1/ni2 ) J, Calculate

the wavelength of light that can excite the electron in a ground state hydrogen atom to n = 7 energy level.


The energy for the transition from n = 1 to n = 7:DE = -2.178 x 10-18 J [1/n2

f - 1/n2i]; nf = 7, ni = 1

DE = -2.178 x 10-18 [1/72 - 1/12] JDE = -2.178 x 10-18 [1/49 - 1/1] JDE = -2.178 x 10-18 [0.02041 - 1] JDE = -2.178 x 10-18 [-0.97959] J = 2.134 x 10-18 J (+, absorption)calculate the l using l = hc/E

6.626 x 10-34 Js x 3.00 x 108 m/s l = ---------------------- 2.13 x 10-18 J l = 9.31 x 10-8 m

Calculation using Bohr eqaution


Wave- Particle Duality of Matter and Energy• Wave theory applies to electromagnetic radiation• EMR can also be described as particles• quanta :A particles of light energy. • Quantum: One particle of light with a certain energy. • Photon: A stream of Quanta• Wave theory could be applied to electrons


Wave theory of the electron• 1924: De Broglie suggested that electrons have

wave properties to account for why their energy was quantized.

• He reasoned that the electron in the hydrogen atom was fixed in the space around the nucleus.

• He felt that the electron would best be represented as a standing wave.

• As a standing wave, each electron’s path must equal a whole number times the wavelength.


De Broglie proposed that all particles have a wavelength as related by:

l = wavelength, metersh = Plank’s constantm = mass, kgv = frequency, m/s

De Broglie waves

l =h

mv


Wave Character of Electrons


What is a wave-mechanical model?• motions of a vibrating string shows one dimensional motion.• Energy of the vibrating string is quantized• Energy of the waves increased with the nodes. • Nodes are places were string is stationary. • Number of nodes gives the quantum number. One

dimensional motion gives one quantum number.Vibrating String : y = sin(npx/l)d2y/dx2 = -(n2p2/l2)sin(npx/l) = -(n2p2/l2)y


Constructively Interfered 2D-Wave


destructively Interfered 2D-Wave


Two-dimensional wave - Vibrations on a Drumskin One circular node

(at the drumskin's edge)

Two circular nodes

(one at the drumskin's edge plus

one more)

Three circular nodes

(one at the drumskin's edge plus

two more)

One transverse node

(plus a circular one at the

drumskin's edge)

Two transverse nodes

(plus one at the drumskin's

edge)


How did Schrodinger come up with a equationstarted with The “Vibrating String” and the "Particle in a One-dimensional Box“ solutionsVibrating String : y = sin(npx/l)d2y/dx2 = -(n2p2/l2)sin(npx/l) = -(n2p2/l2)ySince l = 2l/n; d2y/dx2 = -(4m2v2/h2)y 1/l2= 4/ l2n2; l = h/mv; 1/l2 = 4/ l2n2 = 4m2v2/ h2 Particle in an One-dimensional Box: d2y/dx2 = -(4m2v2p2/h2)y E = ½mv2 + V or v2 = (2/m)(E-V) d2y/dx2 = -(8mp2/h2)(E - V)y


Schrödinger Equation

y = wave function E = total energy V = potential energy


Polar Coordinates


Schrödinger Equation in Polar Coordinates


Solutions to Shrődinger Equation

Series of allowed discrete y values:

yn, l, ml, ms

n = 1,2,3,4,5,6,7..etc.

En = -


Components of yMathematical expression of hydrogen like orbitals

in polar coordinates: y n, l, ml, ms (r,,) = R n, l, (r) Y l, ml, (,)

R n, l, (r ) = Radial Wave Function Y l, ml, (,) =Angular Wave Function


Quantum model of the atom• Schrödinger developed an equation to

describe the behavior and energies of electrons in atoms.

• His equation ( Wave function y ) is similar to one used to describe electromagnetic waves. Each electron can be described in terms of Wave function y its quantum numbers. y n, l, ml, ms),

• y2 is proportional probablity of finding the electron in a given volume. Max Born Interpretation: y2 = atomic orbital


Quantum Model of atom• Electrons travel in three dimensions• Four quantum numbers are needed• three to describe, x, y, z, and four for the spin• four quantum numbers describe an orbital

currently used to explain the arrangement, bonding and spectra of atoms.


Quantum numbers• Principal quantum number, n• Tells the size of an orbital and largely

determines its energy.• n = 1, 2, 3, ……• Angular momentum, l• The number of subshells (s, p, d, f) that a

principal level contains. It tells the shape of the orbitals.

• l = 0 to n - 1


Quantum numbers• Magnetic quantum number, ml

• Describes the direction that the orbital projects in space.

• ml = l to +l (all integers, including zero)• For example, if l = 2, then ml would have

values of -2, -1, 0, 1 and 2.• Knowing all three ml numbers provide us

with a picture of all of the orbitals.


Four Quantum Numbers of the Atom• n value could be 1, 2, 3, 4, 5, 6. 7. . . etc.

• l values depend on n value: can have 0 . . . (n - 1) values

• ml values depends on l value: can have -l . , 0 . . . +l values of ml

• ms values should always be -1/2 or +1/2


Radial Distribution Function, Pnl(r).This is defined as the probability that an electron in

the orbital with quantum numbers n and l will be found at a distance r from the nucleus. It is related to the radial wave function by the following relationship:

R n, l, (r ) = Radial Wave Function Y l, ml, (,) =Angular Wave Function

; normalized by


s orbitals

R n, l, (r) only no Y l, ml, (,)

s-Atomic orbitals


2s

3s

s-Atomic orbitals


p-Atomic orbitals

2p

3p


Nodes in the yTotal nodes = n -1Radial nodes = n -1- lAngular nodes = lEg 4d orbital: Total nodes = 4 -1 = 3Radial nodes = n -1- l = 4-1-2 = 1Angular nodes = l = 2


.

Rnl(r) Pnl(r) n l

1s

1s

1 0

2s

2s

2 0

2p

2p

2 1

3s

3s

3 0

3p

3p

3 1

3d

3d

3 2

Radial wavefunctions, Rnl(r), and the radial distribution functions, Pnl(r)


d-orbitals


Representative d orbitals


f-orbitals


Classification by sublevels

H

Li

Na

Cs

Rb

K

TlHgAuLuBa

Fr

PtIrOsReWTa

He

RnAtPoBiPb

Be

Mg

Sr

Ca

CdAgY PdRhRuTcMoNb

LrRa

ZnCu

Hf

Zr

TiSc NiCoFeMnCrV

In XeITeSbSn

Ga KrBrSeAsGe

Al ArClSPSi

B NeFONC

Gd

Cm

Tb

Bk

Sm

Pu

Eu

Am

Nd

U

Pm

Np

Ce

Th

Pr

Pa

Yb

No

La

Ac

Er

Fm

Tm

Md

Dy

Cf

Ho

Es f

s

d

p


Atomic Orbitals of Multi-Electrnon Atoms• Unlike a hydrogen-like atom multi-electron atoms

there are electron-electron repulsions. • Schrodinger equation cannot be solved

analytically for multi-electron atoms.• However, it is possible to obtain a crude solution

for a multi-electron atom by employing a relatively simple construct.

• The "effective" nuclear charge for each electron is used in place of nuclear charge in the equations for a hydrogen-like atom


Screening (shielding) constant (σ)

a) Screening (shielding) constant (σ) for each electron is calculated based on:

b) the principle quantum numberc) orbital type and penetration and of all other

electrons in an atom.d) σ gives Zeff .

Zeff = Z - σ; Z is the atomic number. σ is the screening constant calculated by Slater Rules


Effective nuclear charge (Zeff)

Zeff is the nuclear charge felt by an electron in a multielectron atom:

a) Each electron in an atom has different Zeff.b) Each Zeff is less than atomic number (Z) since

electrons screen each other from the nucleus.c) Zeff depends on the n and l quantum number of an

electron.d) Zeff Depends on orbital type the electron is in: Zeff

of 4s > 4p > 4d > 4f.


Radial Distribution Functions, Penetration and Shielding


Penetration & Shielding of an Electron in Multi-electron AtomPenetration of an electron:

• Greater the penetration there is more chance of electrons being located close to the nucleus.

• Comparing s, p, d, or f orbitals within same shell (or principle QN), penetration of an electrons are in the order: s > p> d > f

Shielding power of an electron:• Shields of other electrons depends penetration and the

orbital type. Shielding power of electrons in orbitals of that same shell are: s > p > d > f


Slater Rules of Obtaining Zeff

Group electron configuration in the following form:

[1s][2s 2p][3s 3p][3d][4s 4p][4d][4f][5s 5p][5d][5f] etc

Orbitals within a bracket are said to belong to the same group.

• [1s] group where they contribute .30.

• [ns np] group, other electrons in the same group contribute .35

• [ns np] group, each electron in the n-2 or lower group

contributes 1.0.

• [nd] or [nf] group, rules 1 and 2 remain the same and all

electrons in groups to the left contribute 1.0



Consider the outer electron in K. Assume the configuration is [1s2]

[2s2 2p6][3s2 3p6)[3d1] s is then (18 x 1) since the outer electron is in

a [nd] group. Thus Zeff is (19-18)= 1

If we assume that the configuration is [1s2][2s2 2p6][3s2 3p6][3d°][4s1],

the value of s is (8 x 0.85) + (10 x 1)= 16.8 and Zeff is 2.2.

Therefore Zeff is greater and the outer electron experiences more

nuclear attraction when it is in the 4s orbital.



Slater's rule states S = 0.35*x + 0.85*y +z

x,y and z refer to the electron configuration of the atom.

This is for Cl: 1s²2s²2p⁶3s²3p⁵ and for K: 1s²2s²2p⁶3s²3p⁶4s¹

x is the number of valence electrons, the electrons in the highest energy level, 7 for Cl and 1 for K.y is the number of electrons in the energy level below the valence level, 8 for Cl and 8 for K.z is the remaining number of electrons, 2 for Cl and 10 for K.

So we get for Cl S = 0,35*7 + 0,85*8 +2 = 11,25 and for K S =

0,35*1 +0,85*8 + 10 = 17,15


Effective nuclear charge (Zeff) of Atomic Orbitals vs. Z (atomic number)


How do you get the electronic configuration of an atom?

• Use periodic table• Periodic table is divided into orbital blocks• Each period:• represents a shell or n • Start writing electron configuration• Using following order1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d…

(building up (Auf Bau) principle:)


What is Building Up (Auf Bau) Principle• Scheme used by chemist to obtain

electronic configuration of a multi-electron atom in the ground state by filling hydrogen like atomic orbital starting with lowest energy.

• 1s 2s 2p3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s• 5f 6d… (building up principle)• If two or more orbitals exist at the same

energy level, they are degenerate. Do not pair the electrons until you have to.


What is Pauli Exclusion Principle:

Electrons in an atom cannot have all four of their quantum numbers equal.

Eg. He: 1s2

electron orbital n l ml ms

________________________________1s1 1 0 0 +½() 1s2 1 0 0 -½()


Filling order of orbitals


Tl5d

10

6s2

6p1

Hg4f

14

5d10

6s2

Au4f

14

5d10

6s1

Hf4f

14

5d2

6s2

Lu4f

14

5d1

6s2

Li2s

1

Na3s

1

Cs6s

1

Rb5s

1

K4s

1

Fr7s

1

Pt4f

14

5d9

6s1

Ir4f

14

5d7

6s2

Os4f

14

5d6

6s2

Re4f

14

5d5

6s2

W4f

14

5d4

6s2

Ta4f

14

5d3

6s2

H1s

1He1s

2

Rn5d

10

6s2

6p6

At5d

10

6s2

6p5

Po5d

10

6s2

6p4

Bi5d

10

6s2

6p3

Pb5d

10

6s2

6p2

Cd4d

10

5s2

Ag4d

10

5s1

Zr4d

2

5s2

Y4d

1

5s2

Pd4d

10Rh4d

8

5s1

Ru4d

7

5s1

Tc4d

5

5s2

Mo4d

5

5s1

Nb4d

3

5s2

Lr6d

1

7s2

Ba6s

2

Be2s

2

Mg3s

2

Sr5s

2

Ca4s

2

Ra7s

2

Zn3d

10

4s2

Cu3d

10

4s1

Ti3d

2

4s2

Sc3d

1

4s2

Ni3d

8

4s2

Co3d

7

4s2

Fe3d

6

4s2

Mn3d

5

4s2

Cr3d

5

4s1

V3d

3

4s2

In4d

10

5s2

5p1

Xe4d

10

5s2

5p6

I4d

10

5s2

5p5

Te4d

10

5s2

5p4

Sb4d

10

5s2

5p3

Sn4d

10

5s2

5p2

Ga3d

10

4s2

4p1

Kr3d

10

4s2

4p6

Br3d

10

4s2

4p5

Se3d

10

4s2

4p4

As3d

10

4s2

4p3

Ge3d

10

4s2

4p2

Al3s

2

3p1

Ar3s

2

3p6

Cl3s

2

3p5

S3s

2

3p4

P3s

2

3p3

Si3s

2

3p2

B2s

2

2p1

Ne2s

2

2p6

F2s

2

2p5

O2s

2

2p4

N2s

2

2p3

C2s

2

2p2

Gd4f

7

5d1

6s2Cm

5f7

6d1

7s2

Tb4f

9

6s2

Bk5f

9

7s2

Sm4f

6

6s2

Pu5f

6

7s2

Eu4f

7

6s2

Am5f

7

7s2

Nd4f

4

6s2U

5f3

6d1

7s2

Pm4f

5

6s2Np

5f4

6d1

7s2

Ce4f

1

5d1

6s2Th

6f2

7s2

Pr4f

3

6s2Pa

5f2

6d1

7s2

Yb4f

14

6s2

No5f

14

7s2

La5d

1

6s2

Ac6d

1

7s2

Er4f

12

6s2

Fm5f

12

7s2

Tm4f

13

6s2

Md5f

13

7s2

Dy4f

10

6s2

Cf5f

10

7s2

Ho4f

11

6s2

Es5f

11

7s2

Electronic Configuration of elements (core format)


Using the periodic tableTo write the ground-state electron configuration of an element:Starting with hydrogen, go through the elements in order of increasing atomic numberAs you move across a period

• Add electrons to the ns orbital as you pass through groups IA (1) and IIA (2).

• Add electrons to the np orbital as you pass through Groups IIIA (13) to 0 (18).

• Add electrons to (n-1) d orbitals as you pass through IIIB (3) to IIB(12) and add electrons to (n-2) f orbitals as you pass through the f -block.


Writing electron configurations• Examples• O 1s2 2s2 2p4 • Ti 1s2 2s2 2p6 3s2 3p6 3d2 4s2 • Br 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5 • Core format

• O [He] 2s2 2p4 • Ti [Ar] 3d2 4s2 • Br [Ar] 3d10 4s2 4p5


Writing electron configurationsExample - Cl-

• First, write the electron configuration for chlorine:• Cl [Ne] 3s2 3p5

• Because the charge is 1-, add one electron. Cl-

[Ne] 3s2 3p6 or [Ar]


Writing electron configurations

• Electron configurations can also be written for ions.

• Start with the ground-state configuration for the atom.

• For cations, remove a number of the outermost electrons equal to the charge.

• For anions, add a number of outermost electrons equal to the charge.


Writing electron configurationsExample - Ba2+

• First, write the electron configuration for barium.Ba [Xe] 6s2

• Because the charge is 2+, remove two electrons. Ba2+ [Xe] or [Kr] 3d10 4s2 4p6


Hund’s Rule• Rule to fill electrons into p,d,f orbitals containing

more than one sublevel of the same energy. • filling p, d, f orbitals: Put electrons into separate

orbitals of the subshell with parallel spins before pairing electrons.

• The existence of unpaired electrons can be tested for since each acts like a tiny electromagnet.

• Paramagnetic - attracted to magnetic field. Indicates the presence of unpaired electrons.

• Diamagnetic - pushed out of a magnetic field. Indicates that all electrons are paired.


Orbital Box DiagramsValence Shell Electron configuration shown in

box or circle form.


Exception to Building Up Principle a) Electronic Configuration of d-block and f-

block elements d5 or d10 and f7 or f14 are stable Cr :[Ar] 3d4 4s2 wrong Cr :[Ar] 3d5 4s1 correct Cu :[Ar] 3d9 4s2 wrong Cu :[Ar] 3d10 4s1 correct


Lanthanoids

Gd

4f7

5d1

6s2

Tb

4f9

6s2

Sm

4f6

6s2

Eu

4f7

6s2

Nd

4f4

6s2

Pm

4f5

6s2

Ce

4f1

5d1

6s2

Pr

4f3

6s2

Yb

4f14

6s2

La

5d1

6s2

Er

4f12

6s2

Tm

4f13

6s2

Dy

4f10

6s2

Ho

4f11

6s2


Actinoids

Cm

5f7

6d1

7s2

Bk

5f9

7s2

Pu

5f6

7s2

Am

5f7

7s2

U

5f3

6d1

7s2

Np

5f4

6d1

7s2

Th

6f2

7s2

Pa

5f2

6d1

7s2

No

5f14

7s2

Ac

6d1

7s2

Fm

5f12

7s2

Md

5f13

7s2

Cf

5f10

7s2

Es

5f11

7s2


Electronic Configuration of Transition Metal cations

d-block and f-block elements

d orbitals are lower in energy than s orbitals

f orbitals are lower in energy than d orbitals

E.g. Neutral atom Fe :[Ar] 3d6

4s2

Cation, Fe3+

:[Ar] 3d5

Exception to Building Up Principle


Magnetic Properties of Atoms a) Paramagnetism? attracted to magnetic field due to un-paired

electrons. b) Ferromagnetism? attracted very strongly to magnetic field due to

un-paired electrons. c) Diamagnetism? Repelled by a magnetic field due to paired

electrons.


Periodic trends• Many trends in physical and chemical properties

can be explained by electron configuration.• We’ll look at some of the more important

examples.Atomic radiiIonic radiiFirst ionization energiesElectron affinities


How does Zeff vary across a period and down a group?

•Zeff increase going across a period•Zeff decrease going down a group


Types of Atomic Radii

1 Covalent Radii: Radii based on covalently liked atoms in covalently bonded molecules.

2 Van der Waals Radii: Radii based on non bonded atoms in solids.

3 Metallic Radii (12-coordinate):Radii based on metallic solids.

4 Ionic Radii: Radii based on bond distances in ionic solids.


• Atomic radii depend on the distance from the nucleus to the

outermost electron in the valence shell.

• Going across protons are added to nucleus This increase the

Zeff decreasing radii

• Atomic radii decrease going across a period

How does Atomic radii of atoms vary going across a period?


• Atomic radii depend on the distance from the nucleus to the

outermost electron in the valence shell.

• Going down the group outer most shell increases radii hence

the distance from the nucleus

• The atomic radii increase going down a group

How does Atomic radii of elements vary going down a group?


• Cations have smaller radii than neutral atoms.

• Anions have larger radii than neutral atoms

• The more charge on the ion more effect on the

radii.

How does Ionic radii of elements vary?


Atomic radii of elements going down a group?


Atomic radii for the main group (s,p block) elements

Ba

Sr

Ca

Mg

Be

Tl

In

Ga

Al

B

Pb

Sn

Ge

Si

C

Cs

Rb

K

Na

Li

Bi

Sb

As

P

N

Te

Se

S

O

I

Br

Cl

F

H


Atomic radii of the representative- main group elements• Atoms get larger as you go down a group.

A new shell is being added.

• Atoms get smaller as you go across a period.

The nucleus contains more protons.

The higher charge attracts the electrons more strongly, making the atom smaller.


Lanthanoide Contration

• Filling of the 4f orbitals in the lanthanides, which occur within the third series of transition elements, causes these transition metals to be smaller than expected because the 4f orbitals are very poor nuclear shielders and Zeff of 6s2

obitals increase and the atomic radii decrease.• 3rd-series elements have nearly the same

effective nuclear charge as the 2nd-series elements, and thus, nearly the same size

Ce [Xe] 4f1

5d1

6s2


Ionic radii

• Cations• These are smaller than the atoms from

which they are formed.

• For main group elements, the outer shell of electrons is removed.

• The positively charged ion can also do a better job of holding on to the electrons that remain.


Ionic radii

• Anions• These are larger than the atoms from which

there are formed..

• Adding electrons increases the repulsion between electrons.

• The ion has a harder time holding on to the electrons.


Ionic radii (pm)Li Li+ Be Be2+ O O2- F F-

152 74 111 35 74 140 71 133

Na Na+ Mg Mg2+ S S2- Cl Cl-

186 102 160 72 103 184 99 181

K K+ Ca Ca2+ Br Br-

227 138 197 100 114 195

Rb Rb+ Sr Sr2+ I I-

248 149 215 116 133 216

Cs Cs+ Ba Ba2+

265 170 217 136


Isoelectronic configurations

Species that have the same electron configurations.ExampleEach of the following has an electron configuration of 1s2 2s2 2p6

O2- F- Ne

Na+ Mg2+ Al3+


The energy required to remove an electron from an atom.

First Ionization Energy (DH1 ):

Ca ----> Ca+

+ e-; DH1 = positive

Second Ionization Energy (DH2)

Ca+ ----> Ca2+

+ e-; DH2 = positive

DH2 > DH1

What is Ionization Potential?


• Ionization Potential depend on Zeff of the nucleus to the outermost electron

in the valence shell.

• Going down the group Zeff for the outer most shell decrease hence the

Ionization Potential also decrease

• Going across the period Zeff for the outer most shell increase hence the

Ionization Potential also increase

How does Ionization Potential vary going down a group?


Ionization energy

• First ionization energyThe energy to remove one electron from a

neutral atom in the gas phase.• A(g) + first ionization energy A+(g) + e-

• This indicates how easy it is to form a cation. Metals tend to have lower first ionization energies than nonmetals.

• They prefer to become cations.


0

500

1000

1500

2000

2500

0 20 40 60 80 100

First ionization energyHe

Ne

Ar

Kr

Xe

Rn

Firs

t io

niza

tion

ene

rgy

(kJ/m

ol)

Atomic number


Changes of I.E. Across a period


Electron affinity

• A measure of an atom’s tendency to gain electrons in the gas phase.

• A(g) + e- A-(g) + thermal energy

• Electron affinity is an irregular periodic function of atomic number. In general, it increases from left to right.

• Noble gases are not included since they have little or no tendency to gain electrons.


• Electron Affinity depends on Zeff of the nucleus to the outermost electron in the

valence shell.

• Going down the group Zeff for the outer most shell decrease hence the Electron

Affinity also increase

• Going across the period Zeff for the outer most shell increase hence the Electron

Affinity also decrease

How does Electron Affinity vary in the periodic table?


Electron affinity

Atomic number


ElectronegativityPauling Electronegativity, cP

The ability of an atom that is bonded to another atom or atoms to attract electrons to itself.

It is related to ionization energy and electron affinity.

It cannot be directly measured.The values are unitless since they are relative to each

other.The values vary slightly from compound to compound

but still provide useful qualitative predictions.


Electronegativities

0.5

1

1.5

2

2.5

3

3.5

4

0 20 40 60 80 100

Elec

tron

egat

ivit

y

Atomic number

Electronegativity is a

periodic property.


Electronegativity Scales• Pauling Electronegativity, cP

• Mulliken Electronegativity, cM

• The Allred-Rochow, cAR

• Sanderson electronegativity

• Allen electronegativity


Pauling Electronegativity, cP

EA-A and EB-B bond-energy of homonuclear A-A & B-B diatomic moleculesEA-B bond-energy of heteronuclear A-B diatomic moleculecA cB are electronegativity values of A and BPauling comments that it is more accurate to use the geometric mean rather than the arithmetic mean


Mulliken Electronegativity, cM

The Mulliken electronegativity can only be calculated for an element for which the electron affinity is known• For ionization energies and electron affinities in

electronvolts

• For energies in kilojoules per mole


The Allred-Rochow, cAR

The effective nuclear charge, Zeff experienced by valence electrons can be estimated using Slater's rules, while the surface area of an atom in a molecule can be taken to be proportional to the square of the covalent radius, rcov. When rcov is expressed in ångströms,

http://en.wikipedia.org/wiki/Valence_electron

http://en.wikipedia.org/wiki/Slater's_rules

http://en.wikipedia.org/wiki/Covalent_radius

http://en.wikipedia.org/wiki/%C3%85ngstr%C3%B6m


Sanderson, cs

Sanderson has also noted the relationship between electronegativity and atomic size, and has proposed a method of calculation based on the reciprocal of the atomic volume.

The simplest definition of electronegativity is that of Allen, bases on average energy of the valence electrons in a free atom

Allen, cA

where εs,p are the one-electron energies of s-

and p-electrons in the free atom and ns,p are

the number of s- and p-electrons in the valence

shell.

http://en.wikipedia.org/wiki/Valence_electron

Chemistry 281(01) Winter 2014

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Transcript of Chemistry 281(01) Winter 2014