Atomic Structure Inorganic Chem 1

8/10/2019 Atomic Structure Inorganic Chem 1

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Ella KusumastutiKimia Anorganik I

Jurusan Kimia

FMIPA UNNES


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Atomic Structure and Periodic

Table of Elements Perkembangan Teori Atom

Bilangan Kuantum Konfigurasi Elektron (unsur, anion,kation)

Klasifikasi/ penggolongan Unsur dalamSPU

Keperiodikan Sifat Unsur dalam SPU


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What is an atom?

Atom: the smallest unitof matter that retainsthe identity of thesubstance

First proposed byDemocratus

460 BC


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Daltons Atomic Theory

1. All matter is made of tiny indivisibleparticles called atoms.

2. Atoms of the same element areidentical, those of different atoms

are different.3. Atoms of different elements

combine in whole number ratios toform compounds.

4. Chemical reactions involve therearrangement of atoms. No newatoms are created or destroyed.

1808


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Parts of Atoms

J. J. Thomson - English

physicist. 1897

Made a piece of equipment

called a cathode ray tube. It is a vacuum tube - all the air

has been pumped out.

A limited amount of other gasesare put in : Electron

1898

Joseph JohnThompson


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Thomsons Experiment

Voltage source

+-

Metal Disks


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Passing an electric current makes abeam appear to move from the negative

to the positive end

Thomsons Experiment

Voltage source

+-


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Voltage source

Thomsons Experiment

By adding an magnetic field

+

-


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Voltage source

Thomsons Experiment

By adding an magnetic field he foundthat the moving pieces were negative

+

-


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Thomsons Experiment

Used many different metals and gases

Beam was always the same

By the amount it bent he could find theratio of charge to mass

Was the same with every material

Same type of piece in every kind of atom


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Thomsoms Model

Found the electron. Couldnt find

positive (for a while).

Said the atom waslike plum pudding.

A bunch of positive

stuff, with theelectrons able to be

removed.

PLUM PUDDING

MODEL


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Millikans Experiment

Atomizer

Microscope

-

+

Oil

Metal

Plates


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Oil

Atomizer

Microscope

-

+

Oil droplets


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X-rays

X-rays give some drops a charge by knocking

off electrons


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-


+


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They put an electric charge on the plates

++

--


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Some drops would hover

++

--


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+

+ + + + + + +

- - - - - - -

Some drops would hover


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From the mass of the drop and the charge on

the plates, he calculated the charge on an electron

++

--


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Rutherfords Experiment

Ernest Rutherford Englishphysicist. (1910)

Believed the plum pudding modelof the atom was correct.

Wanted to see how big they are. Used radioactivity.

Alpha particles - positivelycharged pieces given off by

uranium. Shot them at gold foil which canbe made a few atoms thick.

1910

ErnestRutherford


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Lead

block

Uranium

Gold Foil

Flourescent

Screen


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He Expected

The alpha particles to pass through

without changing direction very much.

Because

The positive charges were spread out

evenly. Alone they were not enough to

stop the alpha particles.


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What he expected


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Because


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Because, he thought the mass

was evenly distributed in the atom


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Because, he thought

the mass was evenlydistributed in the

atom


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What he got


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How he explained it

+

Atom is mostly empty. Small dense,

positive piece

at center. Alpha particles

are deflected by

it if they get closeenough.


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+


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HISTORY OF THE ATOM

Rutherfords new evidence allowed him to propose a more

detailed model with a central nucleus.

He suggested that the positive chargewas all in a central

nucleus. With this holding the electrons in place by electrical

attraction

However, this was not the end of the story.


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Bohrs Atom Theory

1913 Niels Bohr

studied under Rutherford at the Victoria

University in Manchester.

Bohr refined Rutherford's idea by adding

that the electrons were in orbits. Rather

like planets orbiting the sun. With each

orbit only able to contain a set number of

electrons.


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Bohrs Atom

electrons in orbits

nucleus


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Bohrs Atom


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Bohrs Model

of the Hydrogen Atom(1913)

He proposed that only certain orbits for theelectron are allowed


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Bohrs Empirical Explanation

Electrons can only take discrete energies(energy is related to radius of the orbit)

Electrons can jump between different orbits

due to the absorption or emission of photons

Dark lines in the absorption spectra aredue to photons being absorbed

Bright lines in the emission spectra are

due to photons being emitted

Absorption / Emission of


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Absorption / Emission of

Photons

and Conservation of Energy

Ef-

Ei=

hf

Ei-

Ef=

hf


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Hydrogen Atom is Unstable?

It is known that accelerating charges emit

radiation

Thus, electron should emit radiation, lose energyand eventually fall into the nucleus!

Why doesnt this happen? Shows that somethingwas wrong with this model of the hydrogen atom


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Absorption Spectrum of a Gas

Dark lines will appear in the light spectrum


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Absorption spectrum of

Sun

Emission spectra ofvarious elements


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Balmers Formula for Hydrogen

Notice there are four bright lines in the hydrogen

emission spectrum

Balmer guessed the following formula for thewavelength of these four lines:

where n= 3, 4, 5 and 6


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Energy Levels of Hydrogen


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Electron jumping to

a higher energy level

E = 12.08 eV


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Spectrum of Hydrogen

Bohrs formula:


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Hydrogen atom spectra

Visible lines in H atom

spectrum are called theBALMER series.

High E

Short lHigh n

Low E

Long lLow n

Energy

Ultra Violet

Lyman

Infrared

PaschenVisible

BalmerEn = -1312

n2

65

3

2

1

4

n

Bohrs Quantum Theory of the Atom (1913)


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Bohr s Quantum Theory of the Atom (1913)

Negative electrons move in stable, circular orbits around positive

nuclei

Electrons absorb or emit light by moving out or moving in to other

orbits

Bohr replaced Balmersequations with better ones

Energy levels are far apart at small n, close together at large n

n = 1, 2, 3, etcif the nucleus and electron are completelyseparate

Only worked for H-atom; not a complete description of atomic

structure

22

11

hl

H

nnRE

22

422

)4(

2

h

eZR

o

H

= reduced mass

e = electron charge

Z = nuclear charge

4o= permittivity of vacuum

nucleuse mm

111


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Mechanics Wave Atomic Theory

Subatomic particles (electron, photon, etc) have both

PARTICLE and WAVE properties

Light is electromagnetic radiation - crossed electric

and magnetic waves:

Properties :

Wavelength, l (nm)

Frequency, n (s-1

, Hz)Amplitude, A

constant speed. c

3.00 x 108m.s-1

ELECTROMAGNETIC RADIATION


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Electromagnetic Radiation

wavelength Visible light

wavelengthUltaviolet radiation

Amplitude

Node


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All waves have:

frequency and wavelength symbol: n (Greek letter nu) l (Greek lambda)

units: cycles per sec = Hertz distance (nm)

All radiation: l n = c

where c = velocity of light = 3.00 x 108m/sec


Note: Long wavelengthsmall frequency

Short wavelength

high frequency increasing

wavelengt

increasing

frequency


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Example: Red light has l= 700 nm.

Calculate the frequency, n.

=3.00 x 10

8m/s

7.00 x 10

-7m

4.29 x 1014

Hzn= c

l

Wave nature of light is shown by classical

wave properties such asinterference

diffraction


Q ti ti f E


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Quantization of Energy

Plancks hypothesis:An object can only gain

or lose energy by absorbing or emitting radiant

energy in QUANTA.

Max Planck (1858-1947)

Solved the ultravioletcatastrophe4-HOT_BAR.MOV
http://c/chem_eve/4-HOT_BAR.mov


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E = h n

Quantization of Energy

Energy of radiation is proportional to frequency.

where h = Plancks constant = 6.6262 x 10-34Js

Light with large l(small n) has a small E.

Light with a short l(large n) has a large E.


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Photoelectric effect demonstrates the particle nature of light.

Number of e-ejected does NOT

depend on frequency, rather itdepends on light intensity.

No e-observed until lightof a certain minimum E is used.

Photoelectric Effect

Albert Einstein (1879-1955)


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Photoelectric Effect (2)

Experimental observations can be

explained if light consists of particles

called PHOTONS of discreteenergy.

Classical theory said that E of ejected

electron should increase with increase

in light intensity not observed!

Application of the Schrdinger


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Application of the Schrdinger

Equation to the Hydrogen Atom

The potential energy of the electron-protonsystem is electrostatic:

Use the three-dimensional time-

independent Schrdinger Equation.

S h i l C di t


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Spherical CoordinatesThe potential (central force)V(r)depends on the distance r

between the proton andelectron.

Transform to spherical polar

coordinates because of the radial

symmetry.

The


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The

Schrdinger

Equation inSpherical

CoordinatesTransformed into spherical

coordinates, the

Schrdinger equation

becomes:


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Atomic Line Spectra

Bohrs greatest contribution to

science was in building a simple

model of the atom.

It was based on understanding

the SHARP LINE SPECTRA

of excited atoms.Niels Bohr (1885-1962)(Nobel Prize, 1922)


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Line Spectra of Excited Atoms

Excited atoms emit light of only certain wavelengths The wavelengths of emitted light depend on the

element.

H

Hg

Ne


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Atomic Spectra and Bohr Model

2. But a charged particle moving in anelectric field should emit energy.

+

Electron

orbit

One view of atomic structure in early 20th centurywas that an electron (e-) traveled about the nucleus

in an orbit.

1. Classically any orbit should be

possible and so is any energy.

End result should be destruction!


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Energy of state = - C/n2

whereCis a CONSTANT

n= QUANTUM NUMBER, n = 1, 2, 3, 4, ....

Bohr said classical view is wrong.

Need a new theory now called QUANTUMor

WAVE MECHANICS.

e- can only exist in certain discrete orbits

called stationary states. e- is restricted to QUANTIZEDenergy states.

Atomic Spectra and Bohr Model (2)


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Only orbits where n = integral

number are permitted.

Energy of quantized state = - C/n2

Radius of allowed orbitals

= n2 x (0.0529 nm)

Results can be used toexplain atomic spectra.



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If e-s are in quantized energystates, then DE of states can

have only certain values. This

explains sharp line spectra.

n = 1

n = 2E = -C (1/22)

E = -C (1/12)


H atom

07m07an1.mov

4-H_SPECTRA.MOV

http://c/chem_eve/4-H_SPECTRA.mov


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Calculate DE for e- in H falling fromn = 2 to n = 1 (higher to lower energy) .

n = 1

n = 2

En

ergy

so, E of emitted light = (3/4)R = 2.47 x 1015Hz

and l = c/n = 121.6 nm (in ULTRAVIOLET region)

DE = Efinal- Einitial= -C[(1/12) - (1/2)2] = -(3/4)C

C has been found from experiment. It is now called R,

the Rydberg constant. R = 1312 kJ/mol or 3.29 x 1015Hz

This is exactly in agreement with experiment!

(-ve sign for DE indicates emission (+ve for absorption)

since energy (wavelength, frequency) of light can only be +veit is best to consider such calculations as DE = Eupper- Elower


Hydrogen is therefore a fussy


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Hydrogen is therefore a fussy

absorber / emitter of light

It only absorbs or emits photons with precisely theright energies dictated by energy conservation


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Quantum Numbers and Orbitals

The equations predicted that there arefour quantum numbers.

Principal Quantum Number

n(main energy level or shell)

Angular Quantum Number l(orbital shape)

nl together is called a subshell

Magnetic Quantum Number m(orientation of orbital)

Spin Quantum Number either + or -


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Principal Quantum Number n

Designates the Main Energy Level or Shell anElectron can OccupyOrbital sizes increase as n increases.n2 designates the maximum number of orbitals allowed.2n2designates total electrons in an energy level

n= 1 has only 1 orbital; and 2 electronsn=2 has 4 orbitals; and 8 electrons

n=3 has 9 orbitals; and 18 electrons

A l Q t N b l


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Angular Quantum Number l

Designates the shape of a sublevel l= 0

through (n-1)

The sublevels are

s (sharp) where l=0 p (principal) where l=1

d (diffuse) where l=2

f (fundamental) where l=3

Another name for sublevel is orbital.


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s (sharp) Sublevel

1s

2s

3s

s-orbitals are spherical.There is one s-orbital per shell (n).A total of 2 electrons per s orbital.

No directionality.


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p (principal) Sublevel

P orbitals are peanut shaped.

There are three p-orbitals per shell (n) and have

directionality along the x, y, and z-axis.There are two electrons in each p-orbital.

A total of 6 electrons in all p-orbitals.

Three of these
http://www.uky.edu/~holler/html/p.html


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d (diffuse) Sublevel

d-orbitals are double peanut shaped.

There are five d-orbitals per energy level and havecomplex directionality .

There are 2 electrons per d-orbital.

There are a total of 10 electrons in all d-orbitals.

One of theseTwo of these Two of these
http://www.uky.edu/~holler/html/d.htmlhttp://www.uky.edu/~holler/html/d.html


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f(fundamental) Sublevel

f orbitals are flower shaped.There are seven orbitals and have directionality

There are 2 electrons per f-orbital.There are a total of 14 electrons in all 7 orbitals.

One of these Two of these Two of these Two of these

Angular Quantum Number m
http://www.uky.edu/~holler/html/f.htmlhttp://www.uky.edu/~holler/html/f.htmlhttp://www.uky.edu/~holler/html/f.htmlhttp://www.uky.edu/~holler/html/f.htmlhttp://www.uky.edu/~holler/html/f.html


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Angular Quantum Number m

Designates the orbitals in the subshellOrbitals are oriented on a 3-dimensional axis.

m= -l to +lFor :

l=0 (s); m=0 (-0 to +0)l=1 (p); m=3 (-10+1)l=2 (d); m=5 (-2..-1..0..+1..+2)

l=3 (f); m=7 (-3..-2..-1..0..+1..+2..+3)

There are always 2

electrons per orbital!


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What is a subshell?A subshell is the principal quantum

number n together with the angular

quantum number l.

The n=1 shell has only one subshell which is the 1s subshell.

The n=2 shell has two subshells which are the 2s and 2p subshells.

There are a total of 4 orbitals in these subshells. One in the 2s and

three in the 2p.

Then=3 shell has three subshells which are the 3s, 3p and 3d. ThereAre a total of 9 orbitals in these subshells, one in the 3s, three in the

3p and 5 in the 3d.

Try n=4 for yourself..

Spin Quantum Number + or


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Spin Quantum Number +or -

Designates the spin of each electron in an orbital

Each orbital can hold only 2 electrons.

s has 2e-; p has 6e-; d has 10e-; f has 14e-

2

1

2

1

Electrons like to be in pairs !

Fitting Quantum Numbers Together


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n = # of sublevels per principal energy level n2 = # of orbitals per principal energy level

2n2 = # of electrons per principal energy level

n = 3n = 2n = 1Principallevel (shell)

Sublevel

(subshell)

Orbital

m=-1,0,1 m=-2,-1,0,1,2

Fitting Quantum Numbers Together

s s p s p dl=0 l=1 l=2

m=0

Spin

s= -,+

s py pz dxy dxz dyz dz2 dx2- y2px py pzpx

- +- + - + - + - +- + - + - + - + - +- +- +

Quantum Number Relationships in the


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Quantum Number Relationships in theAtomic Structure

n 1 2 3 4 ...n

l 0 0 1 0 1 2 0 1 2 3

Subshell

designation s s p s p d s p d f

Orbitals in

subshell 1 1 3 1 3 5 1 3 5 7

Subshell

capacity 2 2 6 2 6 10 2 6 10 14

Principal shell

capacity 2 8 18 32 ...2n2


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The Pauli Exclusion Principal

No two electrons can have

the same four quantumnumbers.

O l i O bit l


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All orbitals overlap but electrons cant be more

than 2 er orbital.

Overlapping Orbitals


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Quantum Numbers

ml(magnetic) -l..0..+l Orbital orientation in space

l (angular) 0, 1, 2, .. n-1 Orbital shape or

type (subshell)

n (major) 1, 2, 3, .. Orbital size and energy = -R(1/n2)

Total # of orbitals in lthsubshell = 2 l + 1

Symbol Values Description


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Thank you ....

ATOMIC STRUCTURE


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ATOMIC STRUCTURE

the number of protons in an atom

the number of protons andneutrons in an atomHe

2

4 Atomic mass

Atomic number

number of electrons =number of protons

HELIUM ATOM


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HELIUM ATOM

+

N

N+--

proton

electron neutron

Shell

What do these particles consist of?

ATOMIC STRUCTURE


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ATOMIC STRUCTURE

Electrons are arranged in Energy Levelsor

Shellsaround the nucleus of an atom.

first shell a maximum of 2electrons

second shell a maximum of 8electrons

third shell a maximum of 8electrons

ATOMIC STRUCTURE


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ATOMIC STRUCTURE

There are two ways to represent the atomic

structure of an element or compound;

1. Electronic Configuration

2. Dot & Cross Diagrams

ELECTRONIC CONFIGURATION


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With electronic configuration elements are representednumericallyby the number of electrons in their shells

and number of shells. For example;

N

Nitrogen

7

14

2 in 1stshell

5 in 2ndshell

configuration = 2 , 5

2 + 5 = 7



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Write the electronic configuration for the followingelements;

Ca O

Cl Si

Na20

40

11

23

8

17

16

35

14

28 B 115

a) b) c)

d) e) f)

2,8,8,2 2,8,1

2,8,7 2,8,4 2,3

2,6

DOT & CROSS DIAGRAMS


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With Dot & Cross diagrams elements and compoundsare represented by Dots or Crosses to show electrons,

and circles to show the shells. For example;

Nitrogen N XX X

X

XX

X

N7

14



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Draw the Dot & Cross diagrams for the followingelements;

O Cl8 17

16 35a) b)

O

X

X

X

X

X

X

X

X

Cl

X

X

X

X X

XX

X

X

X

X

X

X

XX

X

X

X

SUMMARY OF ATOMIC STRUCTURE


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1. The Atomic Numberof an atom =number of

protons in the nucleus.

2. The Atomic Massof an atom =number of

Protons + Neutrons in the nucleus.

3. The number of Protons =Number of Electrons.

4. Electrons orbit the nucleus in shells.

5. Each shell can only carry a setnumber of electrons.

Aufbau Approach


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Aufbau Approach

H nds R le


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Hunds Rule

Pauli Exclusion Principle


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Pauli Exclusion Principle

El t i C fi ti


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Electronic Configuration

H atom (1 electron): 1s1

He atom (2 electrons): 1s2

Li atom (3 electrons): 1s2

, 2s1

Cl atom

(17 electrons): 1s2, 2s2, 2p6, 3s2, 3p5

El t i C fi ti


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As atom

33 electons:

1s2

, 2s2

, 2p6

, 3s2

, 3p6

, 4s2

, 3d10

, 4p3

or

[Ar] 4s2, 3d10, 4p3


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Example


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Example

1. 11 Na = 1s 2s 2p 3s

2. 22 Ti = 1s22s22p63s23p64s23d2

n = 3

l = 2 karena orbitalnya dml =

-2 -1 0 +1 +2

Orbitals


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Orbitals

region of probability of finding an

electron around the nucleus

4 types: s, p, d, f

Atomic Orbitals, s-type


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Atomic Orbitals, s type

Atomic Orbitals, p-type


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Atomic Orbitals, d-type


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Mn: [Ar]4s23d?

How many d electrons does Mn have?

4, 5, 6


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Negative ions:

add electron(s), 1 electron for each

negative charge

S-2ion: (16 + 2)electrons:1s2, 2s2, 2p6, 3s2, 3p6



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Positive ions

remove electron(s), 1 electron for each

positive charge

Mg+2ion: (12-2) electrons

1s2, 2s2, 2p6


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How many valence electrons are in Cl,[Ne]3s23p5?

2, 5, 7


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For Cl to achieve a noble gasconfiguration, it is more likely that

electrons would be added

electrons would be removed


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Regions by Electron Type


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Regions by Electron Type


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T d i th P i di T bl


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Trends in the Periodic Table

atomic radius

ionic radius ionization energy

electron affinity

Atomic Radius


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decrease left to right across a period

Zeff= Z - Swhere

Zeff = effective nuclear charge

Z = nuclear charge, atomicnumber

S = shielding constant


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Atomic Radius

Increase top to bottom down a group

Increases from upper right corner tothe lower left corner

Atomic Radius


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Atomic Radius vs. Atomic Number


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Ionic Radii


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I i R di


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Ionic Radius

Same trends as for atomic radius

positive ions smaller than atom negative ions larger than atom

Comparison of Atomic and Ionic Radii


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Ionic Radius


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Ionic Radius

Isoelectronic Series

series of negative ions, noble gas atom,

and positive ions with the same electronic

confiuration

size decreases as positive charge of thenucleus increases

Ionization Energy


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Ionization Energy

energy necessary to remove an electron toform a positive ion

low value for metals, electrons easily

removed

high value for non-metals, electrons

difficult to remove

increases from lower left corner of

periodic table to the upper right corner

Ionization Energies


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Ionization Energies

first ionization energy

energy to remove first electron from an

atom.second ionization energy

energy to remove second electron from a

+1 ion.

etc.

Ionization Energy vs. Atomic Number


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Electron Affinity


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Electron Affinity

energy released when an electron isadded to an atom

same trends as ionization energy,

increases from lower left corner to the

upper right corner

metals have low EA nonmetals have high EA

Magnetism


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g

Result of the spin of electrons

diamagnetism - no unpaired electrons

paramagnetism - one or more unpaired

electrons

Magnetism


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Magnetism

Without applied field With applied field

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Atomic Structure Inorganic Chem 1

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