CHM11-3 Lecture 2
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Transcript of CHM11-3 Lecture 2
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LectureNo. 2(The Electronic Structure of Atoms)
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Content
Atomic Models
Quantum Mechanics
Electronic Configuration
Periodic Relations of Elements
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Atomic Models (Cont.)
John Dalton (1800)
He proposed a modern model based on experimentation.
Daltons Atomic Theory:
1. Elements are composed of extremely small particles called
atoms. All atoms of the same element are alike, and atoms
of different elements are different.
2. The separation of atoms and the union of atoms occur in
chemical reactions. In these reactions, no atom is created or
destroyed, and no atoms of one element are converted intoan atom of another element.
3. A chemical compound is the result of the combination of
atoms of two or more elements. A given compound always
contains the same kinds of atoms combined in the same
proportions.
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Atomic Models (Cont.)
J. J. Thomson (1900)
He discovered electron and proposed a model of the
atom called Plum Pudding Model.
Atoms were made
from a positively
charged substancewith negatively
charged electrons
scattered about.
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Atomic Models (Cont.)
Ernest Rutherford (1910)
He proposed that atoms are mostly empty space and
negative electrons orbit a positive nucleus using the
Gold Foil Experiment.
The Gold Foil Experiment.
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Atomic Models (Cont.)
Rutherfords Model:
A nucleus exists in
the center of the
atom.The nucleus contains
protons and neutrons
which together
account for the mass.
Electrons, which occupy most of the total volume
of the atom, are outside the nucleus and move
rapidly around it.
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Atomic Models (Cont.)
Neils Bohr (1913)
He proposed an improved atomic model.
The Bohrs Model.
Electrons move in
definite orbits aroundthe nucleus.
These orbits, or
energy levels are
located at certaindistances from the
nucleus.
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Atomic Models (Cont.)
Nuclear Bohr Model of Hydrogen Atom:
Electrons normally
exist in the lowerenergy state (ground
state).
When an electron
jumps into higherenergy state it is
said to be in an
exited state.
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Atomic Models (Cont.)
Werner Heisenberg (1925)Theorized that it is impossible to know simultaneously
both the velocity and position of a particles
(Heisenberg Uncertainty Principle).The probable location of an electron is based on how
much energy the electron has.
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Atomic Models (Cont.)
Electron Cloud Model:
Electron cloud is a
space in which electron
are likely to be found.Electrons whirl about
the nucleus billion of
times in 1 second.
They are not movingaround in random
pattern.
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Quantum Mechanics
Electromagnetic Radiation
Light travels through space in a wave motion.
THEWAVENATURE OF LIGHT
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Quantum Mechanics (Cont.)
Definitions
Wavelength () the distance between two
similar points on two successive waves. It
should be expressed in meters. (1nm = 10-9 m).
Amplitude (a) height of a crest or depth of
a trough.
Intensity (brightness) of Radiation
is proportional to the square of amplitude (a2).
Speed of light (c)
equivalent to 2.998 x 108
m/sec.
c =
Frequency () is the
number of waves that pass a
given spot in a second, this is
the reciprocal of second (1/s =hertz).
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Colors of Visible Spectrum:
Color Wavelength, nm
Violet 400 450
Blue 450 500
Green 500 570
Yellow 570 590
Orange 590 620
Red 620 750
Quantum Mechanics (Cont.)
Sample Problem:
What is the frequency of red light with a wavelength of
700nm and violet light with a wavelength of 400nm?
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Quantum Mechanics (Cont.)
Heinrich Hertz
Generated electromagnetic waves with long wavelengths
larger than those of visible light and who demonstrated that
long wavelength radiation exhibits the same phenomena as
light does.
Electromagnetic
Radiation
(Spectra).
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MaxPlank
Proposed the quantum theory of radiant energy.
Suggested that radiant energy could be absorbed orgiven off only in definite quantities called quanta.
E= hv = hc /
hPlanks constant (6.626 x 10-34 J-s)
THEPARTICLENATURE OF LIGHT
Quantum Mechanics (Cont.)
Albert Einstein
Proposed that Planks quanta were discontinuous bits of
energy called photons.
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Quantum Mechanics (Cont.)
Sample Problem:Consider a violet light with a wavelength of 400nm.
Calculate the energy, in joules, of one photon of this
light.
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Isaac Newton
Showed that visible (white light) from the sun can be
broken down into its various component by a prism.
ATOMIC SPECTRA
Quantum Mechanics (Cont.)
Definitions
Continuous Spectrum spreading out
into a wide range band of the white light.
Emission Spectrum When an elementabsorbs sufficient energy (from a flame or
electric arc), it emits radiant energy in the
form of light. When this radiation is passed
through a prism, it separates into a
component wave length.
Absorption Spectrum when
continuous radiation (white light)
passes through a substance, certain
wavelengths of radiation may beabsorbed. These wavelengths are
characteristics of a substance that
absorbs the radiation and pattern of
these lines are referred to as an
absorption spectrum.
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Quantum Mechanics (Cont.)
Every element has its own unique line spectrum,
therefore these spectra are characteristics of an
atoms electronic structure .
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Quantum Mechanics (Cont.)
Bohrs Theory:
1. The electron of the hydrogen atom can exist only in certain
circular orbits (energy levels or shell).
2. The electron has a definite energy characteristic of the orbit
in which it is moving.
3. When an electron of an atom is as close to the nucleus, it isin the condition of the lowest energy called the ground state.
When an atoms are heated in an electric arc or Bunsen
burner, electron absorbs energy and jump to outer levels,
which are higher energy states (excited state).4. When an electron falls back to a lower level, it emits a
definite amount of energy. The energy difference between
the high-energy state and low-energy state is emitted in the
form of a quantum of light.
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Louis De Broglie
He reasoned that if light could show the behavior of
particles (photons) as well as waves, then perhaps an
electron, which Bohr had treated as particle, could behavelike a wave. (Dualistic Nature of Light).
QUANTUM MECHANICAL MODEL
Quantum Mechanics (Cont.)
Erwin Schrdinger
Formulated wave equation that relates the energy of the
electron to its position in the atom.
Solutions of these equations give rise to quantum
numbers.
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ATOMIC ORBITALS
Bohrs Model OUTDATED!!!Schrodingers Equation Quantum Mechanical Model
Heisenberg Uncertainty Principle impossible to know
precisely both the velocity and locations of an electron at
the same time
n = 1 s 1
n = 2 s,p 1,3
n = 3 s,p,d 1,3,5n = 4 s,p,d,f 1,3,5,7
Principal
energy level
Sublevel No. of orbitals
in each sublevel
nucleus
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Quantum Mechanics (Cont.)
Atomic Orbital Shapes:
s and p Orbitals
d Orbitals
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QUANTUM NUMBERS set of numbers that describes an
electron orbital
Quantum Mechanics (Cont.)
1. First Quantum Number (Principal) (n) indicates the
main energy level by the electron. It defines the total energy
of the electrons and has values from (1 to 7).
2. Second Quantum Number (Azimuthal) (l) it describes
the way the electron moves around the nucleus or the shape
of the probability distribution. The values range from 0 to (n-1).
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Quantum Mechanics (Cont.)
n l Spectral Lines1 0 s type (sharp)
2 0
1
s
p type (principal)
3 0
1
2
s
p
d type (diffuse)
4 01
2
3
sp
d
ftype (fundamental)
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Quantum Mechanics (Cont.)
3. Third Quantum Number (Magnetic) (ml) it defines the
possible orientation of the electrons in space. The values arefroml through 0 to +l.
4. Fourth Quantum Number (Spin)
(ms)
it takes intoaccount the spinning of the electron around its own axis as it
moves about the nucleus. The spin is either clockwise or
counterclockwise. The values are +1/2 (clockwise) and -1/2
(counterclockwise).
l ml
Number of orbitals (2l + 1)
0 0 11 -1, 0, +1 3
2 -2,-1, 0, +1, +2 5
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Quantum Mechanics (Cont.)
Aufbaus Principle
Electrons willsuccessively occupy
the available orbitals
on order of
increasing energy.
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Quantum Number
ElectronsAddress
4 #s and no two electrons can have the same quantumnumbers
(n, l, ml, ms)
n = principal energy level
(cannot be zero) n = 1, 2, 3, 4, 5, 6, 7l = sublevel (s, p, d or f) l = n-1
s = 0 p = 1 d = 2 f = 3
ml = orbital ml = -l to +ls _ p _ _ _ d _ _ _ _ _ f _ _ _ _ _ _ _
0 -1 0 +1 -2 -1 0 +1 +2 -3 -2 -1 0 +1 +2 +3
ms
= spin + or -
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Quantum Mechanics (Cont.)
Example:Write the
possible set of
quantum
numbers for theelectrons in:
(a) 3s
(b) 3d
(c) 4f
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Quantum Mechanics (Cont.)
Aufbaus Principle
Electrons willsuccessively occupy
the available orbitals
on order of
increasing energy.
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Hunds Rule
When filling a set of degenerate energy levels, the
electron enter the orbitals singly, with spins in the same
direction (same as s number), until the set is half filled,
before they pair up with opposite spins.
Quantum Mechanics (Cont.)
Concepts:
Paulis Exclusion Principle
Each electron within a given atom must have a unique
set of the four quantum numbers.
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Electronic Configuration
Electronic Configuration
Refers to the arrangement of electrons in energy
levels.Methods ofWriting
1. Orbital method
2. Shell method3.Arrow Rectangular method
4. Core Method
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Electronic Configuration (Cont.)
Example (neutral atom):
Write the electronic configuration using the four
methods of the following elements:
(a) Br
(b) Ca
Example (monoatomic ions):
Give the electronic configuration of:
(a) V4+(b) Cl
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Magnetism
Paramagnetic substances that contain net unpaired
electrons and are attracted by a magnet.Ex. Li
Diamagnetic substances that do not contain net
unpaired electrons and are slightly repelled by amagnet.
Ex. Mg
Quantum Mechanics (Cont.)
RelatedConcepts:
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Periodic Relations ofElements
Periodic Classifications ofElements Elements may be classified according to their
electronic configurations.
1. The Noble Gases
They are also known as InertG
ases orGroup O elements. They are colorless monoatomic gases,
which are chemically unreactive and diamagnetic. They have
outer configurations of ns2np6 (except forHelium).
2. The Representative Elements These elements are foundin the A families of the periodic table. They exhibit a wide
range of chemical behavior and physical characteristics. The
chemistry of these elements depends upon the valence
electrons.
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3. The Transition Elements They are found in the B families
of the periodic table. All of these elements are metals and
most of them are paramagnetic and form highly colored,
paramagnetic compounds.
4. The Inner-Transition Elements These elements are foundat the bottom of the periodic table, but they belong to the 6th
and 7th periods after the elements of group IIIB. All inner-
transition elements are metal and are paramagnetic. Their
compounds are also paramagnetic and colored.
Periodic Relations ofElements (Cont.)
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Periodic Variation in Physical Properties1. Atomic Radius one-half the distance between the nuclei of
the two atoms in an elemental substance.
Periodic Relations ofElements (Cont.)
TREND:
LEFT RIGHT decreases
TOP
BOTTOM increases
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2. Ionic Radius one-half the distance between the nuclei of ametal and a non-metal.
Periodic Relations ofElements (Cont.)
TREND:
LEFT RIGHT decreases*
TOP BOTTOM increases
positive ions are smaller than
the metal atoms from which they
are formed negative ions are larger than
the nonmetal atoms from which
they are formed
* this happens when comparing both
metals and non-metals but by
comparing the metals and non-metals, nonmetals
have larger ionic radius than metals.
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3. Ionization Energy the minimum energy required to remove
an electron from a gaseous atom in its ground state.
Periodic Relations ofElements (Cont.)
TREND:LEFT RIGHT increases
TOP BOTTOM decreases
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Periodic Relations ofElements (Cont.)
4. ElectronAffinity measure of the energy change when
electron is added to a neutral atom to form a negative ion.
TREND:LEFT RIGHT increases
TOP BOTTOM decreases
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END of Lecture No. 2