1. Instructor: Dr. Upali Siriwardane
Chemistry 281(01) Winter 2015
CTH :00-11:15 am
Instructor: Dr. Upali Siriwardane  
Office:  311 Carson Taylor Hall ; Phone: ;
Office Hours:  MTW 8:00 am - 10:00 am;
Th,F 8:30 - 9:30 am & 1:00-2:00 pm.
January 13, Test 1 (Chapters 1&,2),
February 3, 2015 Test 2 (Chapters 2 & 3)
February 26, 2015, Test 3 (Chapters 4 & 5),
Comprehensive Final Make Up Exam: March 3

2. Chapter 1. Atomic Sturcture
Chapter 1.  Atomic structure                                      
Introduction: The origin of the elements                                       3
    The structures of hydrogenic atoms                          
1.1 Spectroscopic information                                           6
1.2 Some principles of quantum mechanics                      8
1.3 Atomic orbitals                                                           
Many-electron atoms                                                   
1.4 Penetration and shielding                                            
1.5 The building-up principle                                            
1.6 The classification of elements
1.7 Atomic properties
                                                       

3. 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).

4. Eight Steps in the History of the Earth
1. The Big Bang
2. Star Formation
3. Supernova Explosion
4. Solar Nebula Condenses
5. Sun & Planetary Rings Form
6. Earth Forms
7. Earth's Core Forms 
8. Oceans & Atmosphere Forms

5. 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 ( MeV/amu)
Binding energy of nuclei expressed in Mev/nucleons

6. Balancing Nuclear Equations

8. Nuclear Binding Energy
The 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: amu
Neutron mass: amu MeV/amu
Electron mass: amu
Mass defect (Dm), then multiply by

9. Bonding Energy Curve

10. 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

11. 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

12. 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 super
nova, spreading the new elements into space with high flux of neutrons to produce heavy elements by neutron capture.

13. Supernova Explosion

14. 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 Supernova
White 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 suddenly
fuses to iron and it explodes known as type 1a supernova.
Most of the elements up to iron in the universe

15. The nucleosynthesis of heavy elements
Havier elements are formed during Supernova explosion.
Giant one star supernova explosions
A 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

16. Cosmic Abundances

17. Terrestrial Abundances

18. Stability of the Elements and Their Isotopes
P/N Ratio
Why are elements
With Z > 82 are
Unstable?

19. Terrestrial Abundances

20. 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.

21. 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.

22. 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.

23. 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 = x [1/n2f - 1/n2i] 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.

24. 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 = x J
En = RH 𝒁𝟐 𝒏𝟐
En = - 𝒎𝒆𝒁𝟐𝒆𝟒 𝟖𝒉𝟐𝟐𝒏𝟐

25. 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 = x J (Z2/n2)
En = x J 1/n2; Z=1 for H

26. Spectra of many other atoms can be described by similar relationships.
Bohr model of the atom
Balmer later determined an empirical relationship that described the spectral lines for hydrogen.
En = - 𝒎𝒆𝒁𝟐𝒆𝟒 𝟖𝒉𝟐𝟐𝒏𝟐
( ) 1
nf2
ni2 - DE
= x J
nf = 2 ni = 3,4, 5, Blamer series
Spectra of many other atoms can be described by
similar relationships.

27. Paschen, Blamer and Lyman Series

28. Calculation using the equation: E = -2
Calculation using the equation: E = x (1/nf2 - 1/ni2 ) J, Calculate the wavelength of light that can excite the electron in a ground state hydrogen atom to n = 7 energy level.

29. Calculation using Bohr eqaution
The energy for the transition from n = 1 to n = 7:
DE = x J [1/n2f - 1/n2i]; nf = 7, ni = 1
DE = x [1/72 - 1/12] J
DE = x [1/49 - 1/1] J
DE = x [ ] J
DE = x [ ] J
= x J (+, absorption)
calculate the l using l = hc/E
6.626 x Js x 3.00 x 108 m/s
l =
2.13 x J
l = x 10-8 m

30. 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

31. 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.

32. De Broglie waves
De Broglie proposed that all particles have a wavelength as related by:
l = wavelength, meters
h = Plank’s constant
m = mass, kg
v = frequency, m/s
l = h mv

33. Wave Character of Electrons

34. 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

35. Constructively Interfered 2D-Wave

36. How did Schrodinger come up with a equation
started with The “Vibrating String” and the "Particle in a One-dimensional Box“ solutions Vibrating String : y = sin(npx/l) d2y/dx2 = -(n2p2/l2)sin(npx/l) = -(n2p2/l2)y Since 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

37.  = wave function E = total energy V = potential energy
Schrödinger Equation
 = wave function E = total energy V = potential energy

38. Polar Coordinates

39. Schrödinger Equation in Polar Coordinates

40. Solutions to Shrődinger Equation
Series of allowed discrete  values:
n, l, ml, ms
n = 1,2,3,4,5,6,7..etc.
En = - 𝒎𝒆𝒁𝟐𝒆𝟒 𝟖𝒉𝟐𝟐𝒏𝟐

41. Components of 
Mathematical expression of hydrogen like orbitals in polar coordinates:
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

42. Quantum model of the atom
Schrödinger developed an equation to describe the behavior and energies of electrons in atoms.
His equation ( Wave function ) is similar to one used to describe electromagnetic waves. Each electron can be described in terms of Wave function its quantum numbers. n, l, ml, ms),
2 is proportional probablity of finding the electron in a given volume. Max Born Interpretation: 2 = atomic orbital

43. 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.

44. 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

45. 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.

46. Four Quantum Numbers of the Atom
n value could be 1, 2, 3, 4, 5, etc.
l values depend on n value: can have (n - 1) values
ml values depends on l value: can have
-l . , l values of ml
ms values should always be -1/2 or +1/2

47. 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                

48. R n, l, (r) only no Y l, ml, (,)
s-Atomic orbitals
R n, l, (r) only no Y l, ml, (,)
s orbitals

49. s-Atomic orbitals 2s 3s

50. p-Atomic orbitals 2p 3p

51. Nodes in the  Total nodes = n -1 Radial nodes = n -1- l
Angular nodes = l
Eg 4d orbital:
Total nodes = = 3
Radial nodes = n -1- l = = 1
Angular nodes = l = 2

52. .
Rnl(r)
Pnl(r) n l 1s
                                                                                       1 2s 2 2p 3s 3 3p 3d
Radial wavefunctions, Rnl(r), and the radial distribution functions, Pnl(r)

53. d-orbitals

54. Representative d orbitals

55. f-orbitals

56. Classification by sublevelsp H He d Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Hf Zr Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Lr La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb f Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

57. 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

58. Screening (shielding) constant (σ)
Screening (shielding) constant (σ) for each electron is calculated based on:
the principle quantum number
orbital type and penetration and of all other electrons in an atom.
σ gives Zeff Zeff = Z - σ; Z is the atomic number σ is the screening constant calculated by Slater Rules

59. Effective nuclear charge (Zeff)
Zeff is the nuclear charge felt by an electron in a multielectron atom:
Each electron in an atom has different Zeff.
Each Zeff is less than atomic number (Z) since electrons screen each other from the nucleus.
Zeff depends on the n and l quantum number of an electron.
Zeff Depends on orbital type the electron is in: Zeff of 4s > 4p > 4d > 4f.

60. Radial Distribution Functions, Penetration and Shielding

61. Penetration & Shielding of an Electron in Multi-electron Atom
Penetration 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

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

63. 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 order
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d… (building up (Auf Bau) principle:)

64. 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.

65. 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 ________________________________
1s ½()
1s ½()

66. 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

67.   Dobereiner Triads  Cl
35.5
Li   7 S  32
Br  79
Na   23
Se  I
127 K 39 Te
128

68. Newlands’ Law of octavesLi Be B C N O F
Octaves 2 Na Mg Al Si P S Cl

69. Lothar Mayer’s atomic volume curves

70. Mendeleyev’s Periodic Table

71. Long Form of Periodic Table

72. 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

73. Chemical properties and the periodic table
Electron 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.

74. 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.

75. Other numbering systems
Previous IUPAC
Current IUPAC and ACS
Preferred US H Li Na K He Be Mg Ca Zn Cu Ti Sc Ni Co Fe Mn Cr V Ga Kr Br Se As Ge Al Ar Cl S P Si B Ne F O N C
IA IIA IIIA IIIA IVA VA VIA VIIA 0
IIIB IVB V B VIB VIIB VIII B IB IIB 1 2 3 4
IA IIA III B IVB VB VIB VIIB VIIIB
IIIA IVA VA VIA VIIA VIIIA IB IIB

76. Filling order of orbitals

77. Filling order of orbitals

78. Electronic Configuration of elements (core format)Tl
5d10
6s2
6p1 Hg
4f14 Au
6s1 Hf
5d2 Lu
5d1 Li
2s1 Na
3s1 Cs Rb
5s1 K
4s1 Fr
7s1 Pt
5d9 Ir
5d7 Os
5d6 Re
5d5 W
5d4 Ta
5d3 H
1s1 He
1s2 Rn
6p6 At
6p5 Po
6p4 Bi
6p3 Pb
6p2 Cd
4d10
5s2 Ag Zr
4d2 Y
4d1 Pd Rh
4d8 Ru
4d7 Tc
4d5 Mo Nb
4d3 Lr
6d1
7s2 Ba Be
2s2 Mg
3s2 Sr Ca
4s2 Ra Zn
3d10 Cu Ti
3d2 Sc
3d1 Ni
3d8 Co
3d7 Fe
3d6 Mn
3d5 Cr V
3d3 In
5p1 Xe
5p6 I
5p5 Te
5p4 Sb
5p3 Sn
5p2 Ga
4p1 Kr
4p6 Br
4p5 Se
4p4 As
4p3 Ge
4p2 Al
3p1 Ar
3p6 Cl
3p5 S
3p4 P
3p3 Si
3p2 B
2p1 Ne
2p6 F
2p5 O
2p4 N
2p3 C
2p2 Gd
4f7
5d1
6s2 Cm
5f7
6d1
7s2 Tb
4f9 Bk
5f9 Sm
4f6 Pu
5f6 Eu Am Nd
4f4 U
5f3 Pm
4f5 Np
5f4 Ce
4f1 Th
6f2 Pr
4f3 Pa
5f2 Yb
4f14 No
5f14 La Ac Er
4f12 Fm
5f12 Tm
4f13 Md
5f13 Dy
4f10 Cf
5f10 Ho
4f11 Es
5f11

79. Using the periodic table
To write the ground-state electron configuration of an element:
Starting with hydrogen, go through the elements in order of increasing atomic number
As 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.

80. 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

81. Writing electron configurations
Example - 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]

82. 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.

83. Writing electron configurations
Example - 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

84. 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.

85. Orbital Box Diagrams
Valence Shell Electron configuration shown in box or circle form.

86. 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

87. La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb
Lanthanoids La
5d1
6s2 Ce
4f1
5d1
6s2 Pr
4f3
6s2 Nd
4f4
6s2 Pm
4f5
6s2 Sm
4f6
6s2 Eu
4f7
6s2 Gd
4f7
5d1
6s2 Tb
4f9
6s2 Dy
4f10
6s2 Ho
4f11
6s2 Er
4f12
6s2 Tm
4f13
6s2 Yb
4f14
6s2

88. Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No
Actinoids Ac
6d1
7s2 Th
6f2
7s2 Pa
5f2
6d1
7s2 U
5f3
6d1
7s2 Np
5f4
6d1
7s2 Pu
5f6
7s2 Am
5f7
7s2 Cm
5f7
6d1
7s2 Bk
5f9
7s2 Cf
5f10
7s2 Es
5f11
7s2 Fm
5f12
7s2 Md
5f13
7s2 No
5f14
7s2

89. Exception to Building Up Principle
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

90. 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.

91. Periodic trends Atomic radii Ionic radii First ionization energies
Many trends in physical and chemical properties can be explained by electron configuration.
We’ll look at some of the more important examples.
Atomic radii
Ionic radii
First ionization energies
Electron affinities

92. How does Zeff vary across a period and down a group?
Zeff increase going across a period
Zeff decrease going down a group

93. 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.

94. 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 across protons are added to nucleus This increase the Zeff decreasing radii
Atomic radii decrease going across a period

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

96. How does Ionic radii of elements vary?
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.

97. Atomic radii of elements going down a group?

98. Atomic radii for the main group (s,p block) elementsLi Be B C N O F Na Mg Al Si P S Cl K Ca Ga Ge As Se Br Rb Sr In Sn Sb Te I Cs Ba Tl Pb Bi

99. 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.

100. 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 5d16s2

101. 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.

102. 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.

103. 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-
K K+ Ca Ca2+ Br Br-
Rb Rb+ Sr Sr2+ I I-
Cs Cs+ Ba Ba2+

104. Isoelectronic configurations
Species that have the same electron configurations.
Example
Each of the following has an electron configuration of 1s2 2s2 2p6
O2- F- Ne
Na+ Mg2+ Al3+

105. What is Ionization Potential?
The energy required to remove an electron from an atom.
First Ionization Energy (DH1 ):
Ca > Ca e-; DH1 = positive
Second Ionization Energy (DH2)
Ca+ ----> Ca e-; DH2 = positive
DH2 > DH1

106. How does Ionization Potential vary going down a group?
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

107. Ionization energy First ionization energy
The 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.

108. First ionization energyHe Ne Ar Kr
First ionization energy (kJ/mol) Xe Rn
Atomic number

109. Changes of I.E. Across a period

110. 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.

111. How does Electron Affinity vary in the periodic table?
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

112. Electron affinity
Atomic number

113. Electronegativity Pauling 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.

114. Electronegativity is a
Electronegativities
Electronegativity is a
periodic property.
Electronegativity
Atomic number

115. Electronegativity Scales
Pauling Electronegativity, cP
Mulliken Electronegativity, cM
The Allred-Rochow, cAR
Sanderson electronegativity
Allen electronegativity

116. Pauling Electronegativity, cP
EA-A and EB-B bond-energy of homonuclear A-A & B-B diatomic molecules
EA-B bond-energy of heteronuclear A-B diatomic molecule
cA cB are electronegativity values of A and B
Pauling comments that it is more accurate to use the geometric mean rather than the arithmetic mean

117. 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

118. 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,

119. 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.