Chemistry 281 01 winter 2014
This presentation is the property of its rightful owner.
Sponsored Links
1 / 124

Chemistry 281(01) Winter 2014 PowerPoint PPT Presentation


  • 47 Views
  • Uploaded on
  • Presentation posted in: General

Chemistry 281(01) Winter 2014. 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.

Download Presentation

Chemistry 281(01) Winter 2014

An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -

Presentation Transcript


Chemistry 281 01 winter 2014

Chemistry 281(01) Winter 2014

CTH 27710:00-11:15 am

Instructor: Dr. UpaliSiriwardane

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.


Chapter 1 atomic sturcture

Chapter 1. Atomic Sturcture

Chapter 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

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

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

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 amumutiplied by the conversion factor (931.5 MeV/amu)

  • Binding energy of nuclei expressed in Mev/nucleons


Balancing nuclear equations

Balancing Nuclear Equations


Nuclear binding energy

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: 1.00728 amu

Neutron mass: 1.00867 amu931.5 MeV/amu

Electron mass: 0.00055 amu

Massdefect (Dm), then multiply by


Bonding energy curve

Bonding Energy Curve


Nuclear fusion reactions

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 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

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 atomic 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.


Nuclear burning

Nuclear Burning


Supernova explosion

Supernova Explosion


The nucleosynthesis of light elements

The nucleosynthesis of light elements

  • Stellar nucleosynthesis

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

  • Double star Supernova

  • White dwarf steals material from another star

  • And get heated huge energy get stored in the while dwarf

  • It goes to nuclear overload and carbon/oxygen

  • Fuse 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

The nucleosynthesis of heavy elements

  • Havier elements are formed during Supernova explosion.

  • Giant one star supernova explosions

  • Heavier star buns all its H and nuclear burning goes faster and forms layer after layers of new elements. Core collapses and become denser.

  • And the star explodes

  • Iron capture neutrons and all heavier elements

  • Corps of supernova explosion leaves a core neutrons. Rotating neutron produces EM pluses creating a pulsar

  • Hypernova explosions: g ray bursts


Cosmic abundances

Cosmic Abundances


Terrestrial abundances

Terrestrial Abundances


Stability of the elements and their isotopes

Stability of the Elements and Their Isotopes

P/N Ratio

Why are elements

With Z > 82 are

Unstable?


Terrestrial abundances1

Terrestrial Abundances


Magic numbers

Magic Numbers

  • Nuclei with either numbers of protons or neutrons equal to Z, N =2, 8, 20, 28, 50, 82, or 126

  • 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

The classification of the elements

  • Dobereiner Triads

  • Newlands called the Law of Octaves

  • Lothar Mayer’s atomic volume curves

  • Mendeleyev’s periodic table


Dobereiner triads

Dobereiner Triads


Newlands law of octaves

Newlands’ Law of octaves


Lothar mayer s atomic volume curves

Lothar Mayer’s atomic volume curves


Mendeleyev s periodic table

Mendeleyev’s Periodic Table


Long form of periodic table

Long Form of Periodic Table


What is periodic table describe its use in chemistry

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.nconn: ns1

Group II- alkaline earth metals:Be, Mg, Ca, Sr, Ba, Ra: Common ele.nconn: 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.nconn ns2 np6


Chemical properties and the periodic table

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

Other ways of numbering groupsin 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

IA IIA III B IVB VB VIB VIIB VIIIB

1 2 13 14 15 16 17 18

IA IIA IIIA IIIA IVA VA VIA VIIA 0

H

He

1

2

3

4

Li

Be

B

C

N

O

F

Ne

IIIA IVA VA VIA VIIA VIIIA IB IIB

3 4 5 6 7 8 9 10 11 12

Na

Mg

Al

Si

P

S

Cl

Ar

IIIB IVB V B VIB VIIB VIII B IB IIB

K

Ca

Sc

Ti

V

Cr

Mn

Fe

Co

Ni

Cu

Zn

Ga

Ge

As

Se

Br

Kr

Other numbering systems

Previous IUPAC

Current IUPAC and ACS

Preferred US


The structures of hydrogenic atoms bohr theory

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

Emission Spectrum of Hydrogen

  • Bohr studied the 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 hydrogen1

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

  • orDE = -2.178 x 10-18 [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.


Bohr model of the atom

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.


What is bohr s atomic model

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 atom1

( )

1

nf2

1

ni2

-

Bohr model of the atom

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

DE

= - 2.178 x 10-18 m-1

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

Spectra of many other atoms can be described by

similar relationships.


Paschen blamer and lyman series

Paschen, Blamer and Lyman Series


Chemistry 281 01 winter 2014

Calculation using the equation: E = -2.178 x 10-18 (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.


Calculation using bohr eqaution

Calculation using Bohr eqaution

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

DE = -2.178 x 10-18 J [1/n2f - 1/n2i]; nf = 7, ni = 1

DE = -2.178 x 10-18 [1/72 - 1/12] J

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

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

DE = -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


Wave particle duality of matter and energy

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

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

h

mv

l =

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


Wave character of electrons

Wave Character of Electrons


What is a wave mechanical model

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.


Constructively interfered 2d wave

Constructively Interfered 2D-Wave


Destructively interfered 2d wave

destructively Interfered 2D-Wave


Two dimensional wave vibrations on a drumskin

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 equation

How did Schrodinger come up with a equation

started with The “Vibrating String” and the "P

article 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 = -(4m2v2p2/h2)y

l = h/mv

Particle in One-dimensional Box:

d2y/dx2 = -(4m2v2p2/h2)y

E = ½mv2 + V or v2 = (2/m)(E-V)

d2y/dx2 = -(8mp2/h2)(E - V)y


Schr dinger equation

Schrödinger Equation

 = wave function

E = total energy

V = potential energy


Polar coordinates

Polar Coordinates


Components of

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


Quantum model of the atom

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


Quantum model of atom

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

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 numbers1

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

Four Quantum Numbers of the Atom

  • nvalue could be 1, 2, 3, 4, 5, 6. 7. . . etc.

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

  • mlvalues 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 p nl r

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:

                     ; normalized by


S atomic orbitals

s-Atomic orbitals

R n, l, (r) only no Y l, ml, (,)

s orbitals


S atomic orbitals1

s-Atomic orbitals

2s

3s


P atomic orbitals

p-Atomic orbitals

2p

3p


Nodes in the

Nodes in the 

Total nodes = n -1

Radial nodes = n -1- l

Angular nodes = l

Eg 4d orbital:

Total nodes = 4 -1 = 3

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

Angular nodes = l = 2


Radial wavefunctions r nl r and the radial distribution functions p nl r

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


D orbitals

d-orbitals


Representative d orbitals

Representative dorbitals


F orbitals

f-orbitals


Classification by sublevels

Ti

Zr

Hf

Classification by sublevels

s

p

H

He

d

Li

Be

B

C

N

O

F

Ne

Na

Mg

Al

Si

P

S

Cl

Ar

K

Ca

Sc

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


Atomic orbitals of multi electrnon atoms

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

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.


Effective nuclear charge z eff

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 nandlquantum number of an electron.

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


Radial distribution functions penetration and shielding

Radial Distribution Functions, Penetration and Shielding


Penetration shielding of an electron in multi electron atom

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 rules of obtaining z eff

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


Slater rules of obtaining z eff1

Slater Rules of Obtaining Zeff

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 rules of obtaining z eff2

Slater Rules of Obtaining Zeff

Slater's rule states S = 0.35*x + 0.85*y +zx,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 z eff of atomic orbitals vs z atomic number

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


How do you get the electronic configuration of an atom

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

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

What is Pauli Exclusion Principle:

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

Eg. He: 1s2electronorbital n lml ms ________________________________

1s1 1 0 0 +½()

1s2 1 0 0 -½()


Filling order of orbitals

Filling order of orbitals


Filling order of orbitals1

Filling order of orbitals


Electronic configuration of elements core format

H

1s1

He

1s2

Li

2s1

Be

2s2

B

2s2

2p1

C

2s2

2p2

N

2s2

2p3

O

2s2

2p4

F

2s2

2p5

Ne

2s2

2p6

Na

3s1

Mg

3s2

Al

3s2

3p1

Si

3s2

3p2

P

3s2

3p3

S

3s2

3p4

Cl

3s2

3p5

Ar

3s2

3p6

K

4s1

Ca

4s2

Sc

3d1

4s2

Ti

3d2

4s2

V

3d3

4s2

Cr

3d5

4s1

Mn

3d5

4s2

Fe

3d6

4s2

Co

3d7

4s2

Ni

3d8

4s2

Cu

3d10

4s1

Zn

3d10

4s2

Ga

3d10

4s2

4p1

Ge

3d10

4s2

4p2

As

3d10

4s2

4p3

Se

3d10

4s2

4p4

Br

3d10

4s2

4p5

Kr

3d10

4s2

4p6

Rb

5s1

Sr

5s2

Y

4d1

5s2

Zr

4d2

5s2

Nb

4d3

5s2

Mo

4d5

5s1

Tc

4d5

5s2

Ru

4d7

5s1

Rh

4d8

5s1

Pd

4d10

Ag

4d10

5s1

Cd

4d10

5s2

In

4d10

5s2

5p1

Sn

4d10

5s2

5p2

Sb

4d10

5s2

5p3

Te

4d10

5s2

5p4

I

4d10

5s2

5p5

Xe

4d10

5s2

5p6

Cs

6s1

Ba

6s2

Lu

4f14

5d1

6s2

Hf

4f14

5d2

6s2

Ta

4f14

5d3

6s2

W

4f14

5d4

6s2

Re

4f14

5d5

6s2

Os

4f14

5d6

6s2

Ir

4f14

5d7

6s2

Pt

4f14

5d9

6s1

Au

4f14

5d10

6s1

Hg

4f14

5d10

6s2

Tl

5d10

6s2

6p1

Pb

5d10

6s2

6p2

Bi

5d10

6s2

6p3

Po

5d10

6s2

6p4

At

5d10

6s2

6p5

Rn

5d10

6s2

6p6

Fr

7s1

Ra

7s2

Lr

6d1

7s2

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

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

Electronic Configuration of elements (core format)


Using the periodic table

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.


Writing electron configurations

Writing electron configurations

  • Examples

  • O1s2 2s2 2p4

  • Ti1s2 2s2 2p6 3s2 3p6 3d2 4s2

  • Br1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5

  • Core format

  • O[He] 2s2 2p4

  • Ti[Ar] 3d2 4s2

  • Br[Ar] 3d10 4s2 4p5


Writing electron configurations1

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 3p6or[Ar]


Writing electron configurations2

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 configurations3

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


Hund s rule

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

Orbital Box Diagrams

Valence Shell Electron configuration shown in box or circle form.


Exception to building up principle

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

Lanthanoids

La

5d1

6s2

Ce

4f1

5d1

6s2

Pr

4f3

6s2

Nd

4f4

6s2

Pm

4f5

6s2

Sm

4f6

6s2

Gd

4f7

5d1

6s2

Tb

4f9

6s2

Dy

4f10

6s2

Ho

4f11

6s2

Er

4f12

6s2

Tm

4f13

6s2

Yb

4f14

6s2

Eu

4f7

6s2


Actinoids

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


Exception to building up principle1

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


Magnetic properties of atoms

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

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

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

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 basesd on bond distances in ionic solids.


How does atomic radii of atoms vary 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 across protons are added to nucleus This increase the Zeff decreasing radii

  • Atomic radii decrease going across a period


How does atomic radii of elements vary going down a group

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

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 of elements going down a group?


Atomic radii for the main group s p block elements

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

Cs

Ba

Tl

Pb

Bi


Atomic radii of the representative main group elements

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

LanthanoideContration

  • 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 Zeffof 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


Ionic radii

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 radii1

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

Ionic radii (pm)

Li Li+Be Be2+OO2-FF-

152 74111 357414071133

Na Na+Mg Mg2+SS2-ClCl-

1861021607210318499181

KK+CaCa2+BrBr-

227138197100114195

RbRb+SrSr2+II-

248149215116133216

CsCs+BaBa2+

265170217136


Isoelectronic configurations

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+


Chemistry 281 01 winter 2014

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+ ----> Ca2+ + e-; DH2 = positive

DH2 > DH1


Chemistry 281 01 winter 2014

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

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

First ionization energy

He

Ne

Ar

Kr

Xe

First ionization energy (kJ/mol)

Rn

Atomic number


Changes of i e across a period

Changes of I.E. Across a period


Electron affinity

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

How does Electron Affinity vary in the periodic table?

  • Electron Affinity depends on Zeffof 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


Electron affinity1

Electron affinity

Atomic number


Electronegativity

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.


Electronegativities

Electronegativities

Electronegativity is a

periodic property.

Electronegativity

Atomic number


Electronegativity scales

Electronegativity Scales

  • Pauling Electronegativity, cP

  • Mulliken Electronegativity, cM

  • The Allred-Rochow, cAR

  • Sanderson electronegativity

  • Allen electronegativity


Pauling electronegativity c p

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

cAcBare 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

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 c ar

The Allred-Rochow, cAR

The effective nuclear charge, Zeffexperienced 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,


Sanderson c s

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.


  • Login