Download Atomic structure

Ella Kusumastuti
Kimia Anorganik I
Jurusan Kimia
FMIPA UNNES
Atomic Structure and Periodic
Table of Elements
• Perkembangan Teori Atom
• Bilangan Kuantum
• Konfigurasi Elektron (unsur, anion,
kation)
• Klasifikasi/ penggolongan Unsur dalam
SPU
• Keperiodikan Sifat Unsur dalam SPU
What is an atom?
460 BC
• Atom: the smallest unit
of matter that retains
the identity of the
substance
• First proposed by
Democratus
Dalton’s Atomic Theory
1. All matter is made of tiny indivisible
2.
3.
4.
1808
particles called atoms.
Atoms of the same element are
identical, those of different atoms
are different.
Atoms of different elements
combine in whole number ratios to
form compounds.
Chemical reactions involve the
rearrangement of atoms. No new
atoms are created or destroyed.
Parts of Atoms
1898
Joseph John
Thompson
• 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 gases
are put in : Electron
Thomson’s Experiment
Voltage source
-
+
Metal Disks
Thomson’s Experiment
Voltage source

+
Passing an electric current makes a
beam appear to move from the negative
to the positive end
Thomson’s Experiment
Voltage source
+
 By adding an magnetic field
Thomson’s Experiment
Voltage source
+
 By adding an magnetic field he found
that the moving pieces were negative
Thomson’s Experiment
• Used many different metals and gases
• Beam was always the same
• By the amount it bent he could find the
ratio of charge to mass
• Was the same with every material
• Same type of piece in every kind of atom
Thomsom’s Model
• Found the electron.
• Couldn’t find
positive (for a while).
• Said the atom was
like plum pudding.
• A bunch of positive
stuff, with the
electrons able to be
removed.
PLUM PUDDING
MODEL
Millikan’s Experiment
Atomizer
+
-
Oil
Microscope
Metal
Plates
Millikan’s Experiment
Atomizer
Oil droplets
+
-
Oil
Microscope
Millikan’s Experiment
X-rays
X-rays give some drops a charge by knocking
off electrons
Millikan’s Experiment
+-
Millikan’s Experiment
-
-
+
+
They put an electric charge on the plates
Millikan’s Experiment
-
-
+
+
Some drops would hover
Millikan’s Experiment
-
-
-
-
-
-
-
+
Some drops would hover
+
+
+
+
+ +
+
Millikan’s Experiment
-
-
+
+
From the mass of the drop and the charge on
the plates, he calculated the charge on an electron
Rutherford’s Experiment
1910
Ernest
Rutherford
• Ernest Rutherford English
physicist. (1910)
• Believed the plum pudding model
of the atom was correct.
• Wanted to see how big they are.
• Used radioactivity.
• Alpha particles - positively
charged pieces given off by
uranium.
• Shot them at gold foil which can
be made a few atoms thick.
Lead
block
Flourescent
Screen
Uranium
Gold Foil
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.
What he expected
Because
Because, he thought the mass
was evenly distributed in the atom
Because, he thought
the mass was evenly
distributed in the
atom
What he got
How he explained it
• Atom is mostly empty.
• Small dense,
positive piece
at center.
• Alpha particles
are deflected by
it if they get close
enough.
+
+
HISTORY OF THE ATOM
Rutherford’s new evidence allowed him to propose a more
detailed model with a central nucleus.
He suggested that the positive charge was all in a central
nucleus. With this holding the electrons in place by electrical
attraction
However, this was not the end of the story.
Bohr’s 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.
Bohr’s Atom
electrons in orbits
nucleus
Bohr’s Atom
Bohr’s Model
of the Hydrogen Atom
(1913)
He proposed that only certain orbits for the
electron are allowed
Bohr’s 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 are
due to photons being absorbed
• Bright lines in the emission spectra are
due to photons being emitted
Absorption / Emission of
Photons
and Conservation of Energy
Ef - Ei = hf
Ei - Ef = hf
Hydrogen Atom is Unstable?
• It is known that accelerating charges emit
radiation
• Thus, electron should emit radiation, lose energy
and eventually fall into the nucleus!
• Why doesn’t this happen? Shows that something
was wrong with this model of the hydrogen atom
Absorption Spectrum of a Gas
Dark lines will appear in the light spectrum
Absorption spectrum of
Sun
Emission spectra of
various elements
Balmer’s Formula for Hydrogen
• Notice there are four bright lines in the hydrogen
emission spectrum
• Balmer guessed the following formula for the
wavelength of these four lines:
where n = 3, 4, 5 and 6
Energy Levels of Hydrogen
Electron jumping to
a higher energy level
E = 12.08 eV
Spectrum of Hydrogen
Bohr’s formula:
Hydrogen atom spectra
High E
Short l
High n
Low E
Long l
Low n
Visible lines in H atom
spectrum are called the
BALMER series.
6
5
4
Energy
3
2
1
En = -1312
n2
Ultra Violet
Lyman
Visible
Balmer
Infrared
Paschen
n
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 Balmer’s equations with better ones
 1
1 
E  RH  2  2 
nh 
 nl
2 2 Z 2 e 4
RH 
(4o ) 2 h 2
1
1
1


 me mnucleus
 = reduced mass
e = electron charge
Z = nuclear charge
4o = permittivity of vacuum
– Energy levels are far apart at small n, close together at large n
n = 1, 2, 3, etc if the nucleus and electron are completely
separate
– Only worked for H-atom; not a complete description of atomic
structure
Mechanics Wave Atomic Theory
ELECTROMAGNETIC RADIATION
• 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 108 m.s-1
Electromagnetic Radiation
wavelength
Visible light
Amplitude
wavelength
Ultaviolet radiation
Node
Electromagnetic Radiation
• All waves have:
frequency
and
• symbol: n (Greek letter “nu”)
wavelength
l (Greek “lambda”)
•
“distance” (nm)
units:
“cycles per sec” = Hertz
• All radiation: l • n = c
where c = velocity of light = 3.00 x 108 m/sec
Note: Long wavelength
 small frequency
Short wavelength
 high frequency
increasing
frequency
increasing
wavelength
Electromagnetic Radiation
Example: Red light has l = 700 nm.
Calculate the frequency, n.
n=
c
l
3.00 x 10
=
8
-7
7.00 x 10
m/s

14
4.29 x 10
Hz
m
• Wave nature of light is shown by classical
wave properties such as
• interference
• diffraction
Quantization of Energy
Max Planck (1858-1947)
Solved the “ultraviolet
catastrophe”
4-HOT_BAR.MOV
• Planck’s hypothesis: An object can only gain
or lose energy by absorbing or emitting radiant
energy in QUANTA.
Quantization of Energy
Energy of radiation is proportional to frequency.
E = h•n
where h = Planck’s constant = 6.6262 x 10-34 J•s
Light with large l (small n) has a small E.
Light with a short l (large n) has a large E.
Photoelectric Effect
Albert Einstein (1879-1955)
Photoelectric effect demonstrates the particle nature of light.
No e- observed until light
of a certain minimum E is used.
Number of e- ejected does NOT
depend on frequency, rather it
depends on light intensity.
Photoelectric Effect (2)
• Classical theory said that E of ejected
electron should increase with increase
in light intensity — not observed!
• Experimental observations can be
explained if light consists of particles
called PHOTONS of discrete
energy.
Application of the Schrödinger
Equation to the Hydrogen Atom
The potential energy of the electron-proton
system is electrostatic:
Use the three-dimensional timeindependent Schrödinger Equation.
Spherical Coordinates
•The potential (central force)
V(r) depends on the distance r
between the proton and
electron.
Transform to spherical polar
coordinates because of the radial
symmetry.
The
Schrödinger
Equation in
Spherical
Coordinates
Transformed into spherical
coordinates, the
Schrödinger equation
becomes:
Atomic Line Spectra
• Bohr’s 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)
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
Atomic Spectra and Bohr Model
One view of atomic structure in early 20th century
was that an electron (e-) traveled about the nucleus
in an orbit.
+
Electron
orbit
1. Classically any orbit should be
possible and so is any energy.
2. But a charged particle moving in an
electric field should emit energy.
End result should be destruction!
Atomic Spectra and Bohr Model (2)
• Bohr said classical view is wrong.
• Need a new theory — now called QUANTUM or
WAVE MECHANICS.
• e- can only exist in certain discrete orbits
— called stationary states.
• e- is restricted to QUANTIZED energy states.
Energy of state = - C/n2
where
C is a CONSTANT
n = QUANTUM NUMBER, n = 1, 2, 3, 4, ....
Atomic Spectra and Bohr Model (3)
Energy of quantized state = - C/n2
• Only orbits where n = integral
number are permitted.
• Radius of allowed orbitals
= n2 x (0.0529 nm)
• Results can be used to
explain atomic spectra.
Atomic Spectra and Bohr Model (4)
If e-’s are in quantized energy
states, then DE of states can
have only certain values. This
explains sharp line spectra.
E = -C
(1/22)
H atom
n=2
07m07an1.mov
E = -C (1/12)
n=1
4-H_SPECTRA.MOV
Calculate DE for e- in H “falling” from
n = 2 to n = 1 (higher to lower energy) .
DE = Efinal - Einitial =
-C[(1/12)
-
(1/2)2]
= -(3/4)C
Energy
Atomic Spectra and Bohr Model (5)
n=2
n=1
• (-ve sign for DE indicates emission (+ve for absorption)
• since energy (wavelength, frequency) of light can only be +ve
it is best to consider such calculations as DE = Eupper - Elower
C has been found from experiment. It is now called R,
the Rydberg constant. R = 1312 kJ/mol or 3.29 x 1015 Hz
so, E of emitted light = (3/4)R = 2.47 x 1015 Hz
and l = c/n = 121.6 nm (in ULTRAVIOLET region)
This is exactly in agreement with experiment!
Hydrogen is therefore a fussy
absorber / emitter of light
It only absorbs or emits photons with precisely the
right energies dictated by energy conservation
Quantum Numbers and Orbitals
“The equations predicted that there are
four quantum numbers.”
 Principal Quantum Number n (main energy level or “shell”)
 Angular Quantum Number  l (orbital shape)
n l together is called a subshell
 Magnetic Quantum Number  m (orientation of orbital)
 Spin Quantum Number  either +½ or -½
Principal Quantum Number “n”
Designates the Main Energy Level or “Shell” an
Electron can Occupy
• Orbital sizes increase as “n” increases.
• n2 designates the maximum number of orbitals allowed.
• 2n2 designates total electrons in an energy level
•n= 1 has only 1 orbital; and 2 electrons
•n=2 has 4 orbitals; and 8 electrons
•n=3 has 9 orbitals; and 18 electrons
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”.
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.
p (principal) Sublevel
Three of these
• 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.
d (diffuse) Sublevel
Two of these
One of these
Two of these
• d-orbitals are “double peanut” shaped.
• There are five d-orbitals per energy level and have
complex directionality .
• There are 2 electrons per d-orbital.
• There are a total of 10 electrons in all d-orbitals.
f (fundamental) Sublevel
One of these
Two of these
Two of these
Two of these
• 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.
Angular Quantum Number “m”
•Designates the “orbitals” in the subshell
•Orbitals are oriented on a 3-dimensional axis.
m= -l to +l
For :
l=0 (s); m=0 (-0 to +0)
l=1 (p); m=3 (-1…0…+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!
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. There
Are 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 “+½ or -½”
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 14eElectrons like to be in “pairs” !
1

2
1

2
Fitting Quantum Numbers Together
n=1
Principal
level (shell)
Sublevel
(subshell)
s
l=0
s
n=3
s
p
p
d
l=1
m=-1,0,1
m=0
s
Orbital
n=2
px
py
l=2
m=-2,-1,0,1,2
pz
px
py
pz
dxy
dxz
dyz
dz2
dx2- y2
-½ +½
-½ +½
-½ +½
-½ +½
-½ +½
s= -½,+½
Spin
• n
-½ +½
-½ +½ -½ +½ -½ +½ -½ +½ -½ +½ -½ +½
= # of sublevels per principal energy level
• n2 = # of orbitals per principal energy level
• 2n2 = # of electrons per principal energy level
Quantum Number Relationships in the
Atomic Structure
n
1
l
0
0
1
0
Subshell
designation
s
s
p
s
p
d
Orbitals in
subshell
1
1
3
5
Subshell
capacity
2
Principal shell
capacity
2
3
1
2
2
3
6
8
4
1
2
2
6
18
10
...n
0
1
2
3
s
p
d
f
1
3
2
5
6
7
10 14
32
...2n2
The Pauli Exclusion Principal
“No two electrons can have
the same four quantum
numbers.”
Overlapping Orbitals
All orbitals overlap but electrons can’t be more
than 2 per orbital.
Quantum Numbers
Symbol
Values
Description
n (major)
1, 2, 3, ..
Orbital size and energy = -R(1/n2)
l (angular)
0, 1, 2, .. n-1
Orbital shape or
type (subshell)
ml (magnetic)
-l..0..+l
Orbital orientation in space
Total # of orbitals in lth subshell = 2 l + 1
Thank you ....