chem281-chapter1-w15 - Louisiana Tech University

Chemistry 281(01) Winter 2015

CTH 277 10:00-11:15 am

Instructor: Dr. Upali Siriwardane

E-mail : upali@latech.edu

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

4941;

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

Th,F 8:30 - 9:30 am & 1:00-2:00 pm.

January 13, 2015 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

Chemistry 281, Winter 2015, LA Tech Chapter-1-1

Chapter 1. Atomic Sturcture

Chapter 1. Atomic structure

Introduction: The origin of the elements

The structures of hydrogenic atoms

1.1 Spectroscopic information

1.2 Some principles of quantum mechanics

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

9

15

15

17

3

4

6

8

20

22


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 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


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=mc 2

•

• Nuclear reactions are balanced by A (mass) and

Z (atomic) number.

Energy released is E=mc 2 , 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 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. (

D m) –mass defect

(

D m) = Mass of Nuclide - mass of (p + n +e )

Proton mass : 1.00728 amu

Neutron mass : 1.00867 amu 931.5 MeV/amu

Electron mass : 0.00055 amu

Mass defect (

D m), 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 ( 2 H), 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 super nova, spreading the new elements into space with high flux of neutrons to produce heavy elements by neutron capture.


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 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


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


Cosmic Abundances


Terrestrial Abundances


Stability of the Elements and Their

Isotopes

P/N Ratio

Why are elements

With Z > 82 are

Unstable?


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 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:

D

E = E f

• or

D

E = -2.178 x 10 -18 [1/n 2 f

- 1/n 2 i

] J

- E i

• D

E is negative for an emission and positive for an absorption

• D

E 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.

• R

H

= 2.178 x 10 -18 J

E n

= R

H

𝒁 𝟐 𝒏 𝟐

E n

=

𝒎 𝒆

𝒁 𝟐 𝒆 𝟒

𝟖𝒉

𝟐





𝟐 𝒏

𝟐


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 (E n

):

• E n

• E n

= -2.178 x 10

= -2.178 x 10

-18

-18

J (Z 2

J 1/n

/n

2 ;

2 )

Z=1 for H


Bohr model of the atom

Balmer later determined an empirical relationship that described the spectral lines for hydrogen.

E n

=

𝒎 𝒆

𝒁 𝟐 𝒆 𝟒

𝟖𝒉 𝟐





𝟐 𝒏 𝟐

D E = - 2.178 x 10

-18

J

1 ( ) f

2 n

1

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

Spectra of many other atoms can be described by similar relationships.


Paschen, Blamer and Lyman Series


Calculation using the equation:

E = -2.178 x 10 -18 (1/n f

2 - 1/n i

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


Calculation using Bohr eqaution

The energy for the transition from n = 1 to n = 7:

D

E = -2.178 x 10 -18 J [1/n 2 f

- 1/n

D

E = -2.178 x 10 -18 [1/7 2 - 1/1 2 ] J

2 i

]; n f

D

E = -2.178 x 10 -18 [1/49 - 1/1] J

= 7, n i

= 1

D

E = -2.178 x 10 -18 [0.02041 - 1] J

D

E = -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 10 8 m/s l

= ----------------------

2.13 x 10 -18 J l

= 9.31 x 10 -8 m


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 waves

De Broglie proposed that all particles have a wavelength as related by: l = h mv l

= wavelength, meters h = Plank’s constant m = mass, kg v = frequency, m/s


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(n p x/l) d 2 y

/dx 2 = -(n 2 p

2 /l 2 )sin(n p x/l) = -(n 2 p

2 /l 2 ) y


Constructively Interfered 2D-Wave


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(n p x/l) d 2 y

/dx 2 = -(n 2 p

2 /l 2 )sin(n p x/l) = -(n 2 p

2 /l 2 ) y

Since l

= 2l/n; d 2 y

/dx 2 = -(4m 2 v 2 /h 2 ) y

1/ l 2 = 4/ l

2 n 2 ; l

= h/mv; 1/l 2 = 4/ l

2 n 2 = 4m 2 v 2 / h 2

Particle in an One-dimensional Box : d 2 y

/dx 2 = -(4m 2 v 2 p

2 /h 2 ) y

E = ½mv 2 + V or v 2 = (2/m)(E-V) d 2 y

/dx 2 = -(8m p

2 /h 2 )(E - V) y


Schrödinger Equation y

= wave function E = total energy V = potential energy


Polar Coordinates


Schrödinger Equation in Polar Coordinates


Solutions to Shr ődinger Equation

Series of allowed discrete y values: y n, l, ml, ms n = 1,2,3,4,5,6,7..etc.

E n

=

𝒎 𝒆

𝒁 𝟐 𝒆 𝟒

𝟖𝒉 𝟐





𝟐 𝒏 𝟐


Components of y

Mathematical 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, m l

, m s

), y 2 is proportional probablity of finding the electron in a given volume. Max Born

Interpretation : y 2

= 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, m l

Describes the direction that the orbital projects in space.

•

•

• m l

= l to + l (all integers, including zero)

For example, if l = 2, then m l values of -2, -1, 0, 1 and 2.

would have

Knowing all three m l 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

• m l values depends on l value: can have

-l . , 0 . . . +l values of m l

• m s values should always be -1/2 or +1/2


Radial Distribution Function, P nl

(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-Atomic orbitals

R n, l

, (r) only no Y l

, m l

, (



,



)

Chemistry 281, Winter 2015, LA Tech s orbitals

Chapter-1-48

s-Atomic orbitals

2s

3s


p-Atomic orbitals

2p

3p


Nodes in the y

Total nodes = n -1

Radial nodes = n -1l

Angular nodes = l

Eg 4d orbital:

Total nodes = 4 -1 = 3

Radial nodes = n -1l = 4-1-2 = 1

Angular nodes = l = 2


.

Radial wavefunctions, R nl

(r), and the radial distribution functions, P nl

(r)

R nl

(r) P nl

(r) n l

1s 1s 1 0

2s 2s 2 0

2p

3s

3p

3d

Chemistry 281, Winter 2015, LA Tech

2p

3s

3p

3d

2

3

3

1

0

1

3 2

Chapter-1-52

d-orbitals


Representative d orbitals


f-orbitals


Classification by sublevels s

H

Li Be

Na Mg d p

He

B C N O F Ne

Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Lr f


La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb

Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No

Chapter-1-56

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) b) c) d)

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 Z eff

Z eff

.

= Z σ; Z is the atomic number .

σ is the screening constant calculated by Slater

Rules


Effective nuclear charge (Z eff

)

Z eff is the nuclear charge felt by an electron in a multielectron atom: a) b) c) d)

Each electron in an atom has different Z eff

.

Each Z eff is less than atomic number (Z) since electrons screen each other from the nucleus.

Z eff depends on the electron.

n and l quantum number of an

Z eff

Depends on orbital type the electron is in: Z of 4s > 4p > 4d > 4f.

eff


Radial Distribution Functions, Penetration and Shielding


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


Slater Calculation of (Z eff

)


Slater Calculation of (Z eff

)


11. Cu: (1s 2 )(2s 2 , 2p 6 ) (3s 2 ,3p 6 ) (3d 10 ) (4s 1 ) : there are two possible scenarios for forming Cu+ ionionizing 3d 10 electron or 4s 1 . Using Slater’s Rules show which one of the electrons 4s or 3d would come out easily.

If the electron is in a d or f-orbital :

All electrons in groups higher than the electron in question contribute zero to s

.

Each electron in the same group contributes 0.35 to s

.

All those in groups to the left contribute 1.0 to s

(n-3) (n-2) (n-1) (n-1)

Cu: (1s 2 )(2s 2 , 2p 6 ) (3s 2 ,3p 6 ) (3d 10 ) (4s 1 ) s

(4s 1 ) = ( 10x1 ) ( 8x 0.85)(1X10) = 26.8 Z eff

Cu: (1s 2 )(2s 2 , 2p 6 ) (3s 2 ,3p 6 ) (3d 10 ) (4s 1 ) s

(3d 1 ) = ( 18x1 ) ( 8x 0.35) (0) = 20.8 Z eff

= 29 – 26.6 = 2.4

= 29 – 20.8 = 8.2


Effective nuclear charge (Z eff

) 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 order

1s 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: 1s 2 electron orbital n l m l m s

________________________________

1s 1 1 0 0 +½(  )

1s 2 1 0 0 ½(  )


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 con n : ns 1

Group II- alkaline earth metals :Be, Mg, Ca, Sr, Ba,

Ra: Common ele.

n con n : ns 2

Group VII- Halogens : Cl, Br, I, At:

Common ele.

n con n :ns 2 np 5

Group VIII- Noble gases :He, Ne, Ar, Kr, Xe, Rn:

Common ele.

n con n ns 2 np 6


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.


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

Previous IUPAC

Current IUPAC and ACS

Preferred US

IA IIA

1 2

IA IIA

H

3

4

1

2

Li

Na

Be

Mg

IIIA IVA VA VIA VIIA VIIIA

3 4 5 6 7 8 9 10 11 12

IIIB IVB V B VIB VIIB VIII B IB IIB

III B IVB VB VIB VIIB VIIIB

13 14 15 16 17 18

IIIA IIIA IVA VA VIA VIIA 0

Al

C

Si

N

P

O

S

F

Cl

He

Ne

Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr


Filling order of orbitals


Filling order of orbitals


Electronic Configuration of elements (core format)

Li

2s

1

Be

2s

2

B

2s

2

2p

1

C

2s

2p

2

2

N

2s

2p

2

3

O

2s

2p

2

4

F

2s

2p

2

5

Ne

2s

2p

2

6

H

1s

1

He

1s

2

Na

3s

1

Mg

3s

2

Al

3s

2

3p

1

Si

3s

3p

2

2

P

3s

3p

2

3

S

3s

3p

2

4

Cl

3s

3p

2

5

Ar

3s

3p

2

6

K

4s

1

Ca

4s

2

Sc

3d

4s

2

Ti

3d

4s

2

2

V

3d

3

4s

2

Rb

5s

1

Cs

6s

Sr

5s

2

Ba

6s

Y

1

4d

5s

2

Zr

4d

5s

2

2

Lu

4f

14

5d

6s

1

2

Hf

4f

14

5d

6s

2

2

Fr

7s

1

Ra

7s

2

Lr

1

6d

7s

2


La

1

5d

6s

2

Ac

6d

7s

2

Nb

4d

3

5s

2

Ta

4f

5d

6s

3

2

Cr

3d

5

4s

1

Mo

4d

5

5s

1

W

4f

5d

6s

4

2

Ce

4f

5d

6s

1

1

2

Th

6f

2

7s

2

Pr

4f

3

6s

2

Pa

5f

6d

7s

1

2

Mn

3d

5

4s

2

Tc

4d

5

5s

2

Re

4f

5d

6s

5

2

Fe

3d

6

4s

2

Ru

4d

7

5s

1

Os

4f

5d

6s

6

2

Co

3d

7

4s

2

Rh

4d

8

5s

1

4f

Ir

5d

6s

7

2

Ni

8

3d

4s

2

Pt

4f

14

5d

6s

9

1

Cu

3d

4s

1

Pd

4d

10

Ag

4d

5s

1

Zn

3d

10

4s

2

Cd

4d

10

5s

2

Au

4f

5d

14

6s

10

1

Hg

4f

5d

6s

10

2

Ga

3d

4s

4p

2

1

In

4d

5s

5p

2

1

Tl

5d

6s

6p

2

1

Nd

4f

4

6s

2

U

5f

6d

7s

3

1

2

Pm

4f

5

6s

2

Np

5f

4

6d

7s

1

2

Sm

4f

6

6s

2

Pu

5f

6

7s

2

Eu

4f

7

6s

2

Am

5f

7

7s

2

Gd

4f

5d

6s

7

1

2

Cm

5f

6d

7s

7

1

2

Tb

4f

9

6s

2

Bk

5f

9

7s

2

Dy

4f

6s

2

Cf

5f

7s

2

Ge

3d

4s

4p

10

2

2

Sn

4d

5s

5p

10

2

2

Pb

5d

6s

6p

10

2

2

As

3d

4s

4p

10

2

3

Sb

4d

5s

5p

10

2

3

Bi

5d

6s

6p

10

2

3

Se

3d

4s

4p

10

2

4

Te

4d

5s

5p

10

2

4

Po

5d

6s

6p

10

2

4

Br

3d

4s

4p

10

2

5

4d

5s

5p

I

10

2

5

At

5d

6s

6p

10

2

5

Kr

3d

4s

4p

10

2

6

Xe

4d

5s

5p

10

2

6

Rn

5d

6s

6p

10

2

6

Ho

4f

6s

2

Es

5f

7s

2

Er

4f

12

6s

2

Tm

4f

6s

2

Yb

4f

6s

2

Fm

5f

7s

2

Md

5f

7s

2

No

5f

7s

2

Chapter-1-81

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 n s orbital as you pass through groups

IA (1) and IIA (2).

• Add electrons to the n p 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

Ti

Br

Core format

1 s 2 2 s 2 2 p 4

1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 2 4 s 2

1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 3 d 10 4 s 2 4 p 5

O

Ti

Br

[He] 2 s 2 2 p 4

[Ar] 3 d 2 4 s 2

[Ar] 3 d 10 4 s 2 4 p 5



•

•

Example - Cl -

First, write the electron configuration for chlorine:

Cl [Ne] 3 s 2 3 p 5

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

[Ne] 3 s 2 3 p 6 or [Ar]



•

•

•

•

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.



•

•

Example - Ba 2+

First, write the electron configuration for barium.

Ba [Xe] 6 s 2

Because the charge is 2+, remove two electrons.

Ba 2+ [Xe] or [Kr] 3 d 10 4 s 2 4 p 6


•

•

•

•

•

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 Diagrams

Valence Shell Electron configuration shown in box or circle form.


Exception to Building Up Principle a) Electronic Configuration of d-block and fblock elements d 5 or d 10 and f 7 or f 14 are stable

Cr :[Ar] 3d 4 4s 2 wrong

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

Cu :[Ar] 3d 9 4s 2 wrong

Cu :[Ar] 3d 10 4s 1 correct


Lanthanoids

La

5d

1

6s

2

Ce

4f

1

5d

1

6s

2

Pr

4f

6s

3

2

Nd

4f

4

6s

2

Pm

4f

5

6s

2

Sm

4f

6s

6

2

Eu

4f

7

2

6s

Gd

4f

7

5d

1

6s

2

Tb

4f

9

6s

2

Dy Ho

4f

10

6s

2

4f

11

6s

2

Er Tm Yb

12

4f

6s

2

4f

13

6s

2

4f

14

6s

2


Actinoids

Ac

6d

1

7s

2

Th

6f

2

7s

2

Pa

2

5f

1

6d

2

7s

U

5f

3

1

6d

7s

2

Np

5f

4

6d

1

7s

2

Pu

5f

6

2

7s

Am

Cm

5f

7

2

7s

5f

7

6d

1

7s

2

Bk

5f

9

2

7s

Cf Es

5f

10

2

7s

5f

11

7s

2

Fm Md No

12

5f

7s

2

5f

13

2

7s

5f

14

7s

2


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] 3d

6

4s

2

Cation, Fe

3+

:[Ar] 3d

5


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 radii

Ionic radii

First ionization energies

Electron affinities


How does Z eff vary across a period and down a group?

• Z eff

• Z eff increase going across a period 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.


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 Z eff decreasing radii

• Atomic radii decrease going across a period


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


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.


Atomic radii of elements going down a group?


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

H

Li 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

Bi Cs Ba


Tl Pb

Chapter-1-101

Atomic radii of the representativemain 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 Z eff of 6 s obital increase and the atomic radii decrease.

2

3rd-series d-elements have nearly the same effective nuclear charge as the 2nd-series delements, and thus, nearly the same size

Ce: [Xe] 4f

1

5d

1

6s

2


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

152 74

+ Be Be 2+

111 35

Na Na + Mg Mg

186 102 160 72

2+

K K + Ca Ca 2+

227 138 197 100

Rb Rb + Sr Sr 2+

248 149 215 116

Cs Cs + Ba Ba 2+

265 170 217 136


O

74

O 2F

140 71

F -

133

S S 2Cl

103 184 99

Cl -

181

Br Br -

114 195

I I -

133 216

Chapter-1-106

Isoelectronic configurations

Species that have the same electron configurations.

Example

Each of the following has an electron configuration of 1 s 2 2 s 2 2 p 6

O 2F Ne

Na + Mg 2+ Al 3+


What is Ionization Potential?

The energy required to remove an electron from an atom.

First Ionization Energy (

D

H1 ):

Ca ----> Ca + + e-;

D

H1 = positive

Second Ionization Energy (

D

H2)

Ca+ ----> Ca 2+ + e-;

D

H2 = positive

D

H2 >

D

H1


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


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.


First ionization energy

2500

He

Ne

2000

Ar

1500

Kr

Xe

1000

500

0

0


20 40

Atomic number

60

Rn

80 100

Chapter-1-111

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.


How does Electron Affinity vary in the periodic table?

• Electron Affinity depends on Z eff of the nucleus to the outermost electron in the valence shell.

• Going down the group Z eff for the outer most shell decrease hence the Electron Affinity also increase

• Going across the period Z eff for the outer most shell increase hence the Electron Affinity also decrease


Electron affinity


Atomic number

Chapter-1-115

Electronegativity

Pauling Electronegativity, c

P

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

4

3.5

3

2.5

2

1.5

1

0.5

0 20


40

Atomic number

60

Electronegativity is a periodic property.

80 100

Chapter-1-117

•

Electronegativity Scales


P

• Mulliken Electronegativity, c

M

• The Allred-Rochow, c

AR

• Sanderson electronegativity

• Allen electronegativity



P

E

A-A and

E

B-B bond-energy of homonuclear A-A & B-B diatomic molecules

E

A-B bond-energy of heteronuclear A-B diatomic molecule c

A c

B are electronegativity values of A and B

Pauling comments that it is more accurate to use the geometric mean rather than the arithmetic mean


Mulliken Electronegativity, c

M

•

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, c

AR

The effective nuclear charge, Z eff 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 , r cov

. When r cov expressed in ångströms , is


Sanderson, c s

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.

Allen, c

A

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

Chemistry 281, Winter 2015, LA Tech where ε s,p are the oneelectron energies of s- and p-electrons in the free atom and n s,p are the number of s- and p-electrons in the valence shell.

Chapter-1-122

chem281-chapter1-w15 - Louisiana Tech University

Bonding Energy Curve

Supernova Explosion

Terrestrial Abundances

Terrestrial Abundances

Bohr model of the atom

Wave- Particle Duality of Matter and Energy

Wave theory of the electron

Four Quantum Numbers of the Atom

p-Atomic orbitals

The classification of the elements

Newlands’ Law of octaves

What is periodic table?

Describe its use in chemistry?

Chemical properties and the periodic table

Writing electron configurations

Magnetic Properties of Atoms

Periodic trends

Types of Atomic Radii

Lanthanoide Contration

First ionization energy

Changes of I.E. Across a period

Electron affinity

Electronegativity

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chem281-chapter1-w15 - Louisiana Tech University

Bonding Energy Curve

Supernova Explosion

Terrestrial Abundances

Terrestrial Abundances

Bohr model of the atom

Wave- Particle Duality of Matter and Energy

Wave theory of the electron

Four Quantum Numbers of the Atom

p-Atomic orbitals

The classification of the elements

Newlands’ Law of octaves

What is periodic table?

Describe its use in chemistry?

Chemical properties and the periodic table

Writing electron configurations

Magnetic Properties of Atoms

Periodic trends

Types of Atomic Radii

Lanthanoide Contration

First ionization energy

Changes of I.E. Across a period

Electron affinity

Electronegativity

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