Atomic Structure Inorganic Chem 1
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Transcript of 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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Millikans Experiment
Oil
Atomizer
Microscope
-
+
Oil droplets
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Millikans Experiment
X-rays
X-rays give some drops a charge by knocking
off electrons
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-
Millikans Experiment
+
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Millikans Experiment
They put an electric charge on the plates
++
--
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Millikans Experiment
Some drops would hover
++
--
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Millikans Experiment
+
+ + + + + + +
- - - - - - -
Some drops would hover
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Millikans Experiment
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
Electromagnetic Radiation
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
Electromagnetic Radiation
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
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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.
Atomic Spectra and Bohr Model (3)
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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)
Atomic Spectra and Bohr Model (4)
H atom
07m07an1.mov
4-H_SPECTRA.MOV
Atomic Spectra and Bohr Model (5)
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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
Atomic Spectra and Bohr Model (5)
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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ELECTRONIC CONFIGURATION
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
ELECTRONIC CONFIGURATION
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ELECTRONIC CONFIGURATION
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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DOT & CROSS DIAGRAMS
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
DOT & CROSS DIAGRAMS
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DOT & CROSS DIAGRAMS
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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Electronic Configuration
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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Electronic Configuration
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Electronic Configuration
Negative ions:
add electron(s), 1 electron for each
negative charge
S-2ion: (16 + 2)electrons:1s2, 2s2, 2p6, 3s2, 3p6
Electronic Configuration
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Electronic Configuration
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